What are Catalysts? Revision Notes | A Level AQA Chemistry

1. Why Catalysts Matter

Picture the Haber process: nitrogen and hydrogen sitting in a reactor at high pressure, barely reacting at room temperature. Add iron, and the reaction rate jumps dramatically. That iron is a catalyst — and understanding exactly why it works is the heart of this topic.

Catalysts are not a chemical shortcut or a trick. They operate through a precise kinetic mechanism that the AQA specification requires you to explain at the molecular level. By the end of this lesson, you will be able to draw on the Maxwell-Boltzmann distribution to justify why a lower activation energy translates directly into a faster reaction rate.

2. Defining a Catalyst: Getting the Wording Right

The AQA definition is precise, and every word in it earns marks. A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount at the end of the reaction.

Two parts of that definition are frequently tested separately. First, the catalyst is not changed in chemical composition — it does not become a different substance. Second, it is not changed in amount — the same number of moles is present before and after. A catalyst can be temporarily consumed in an intermediate step, but it is always regenerated by the end of the overall reaction.

3. How Catalysts Work: The Alternative Pathway

Here is the core mechanism. Every reaction has an activation energy — the minimum energy that colliding particles must possess for the reaction to occur. Without a catalyst, reactant molecules must climb over a high energy barrier to reach the transition state and form products.

The height difference between start and destination is unchanged, just as ΔH is unchanged by a catalyst: the catalyst only lowers the activation energy barrier.

Notice three things in that profile. The reactants start at the same energy level in both pathways. The products end at the same energy level — confirming $\Delta H$ is unchanged. Only the peak of the energy barrier, the activation energy, is lower in the catalysed route.

4. The Maxwell-Boltzmann Distribution and Catalysis

The reaction profile tells us the activation energy is lower. The Maxwell-Boltzmann distribution tells us why that matters so much for reaction rate. This is where the two ideas connect, and it is exactly what AQA asks you to explain.

The Maxwell-Boltzmann distribution is a graph showing how molecular energies are spread across a sample of gas at a fixed temperature. The x-axis shows molecular energy, the y-axis shows the number of molecules with that energy. The curve rises from zero, peaks, then tails off asymptotically — it never quite touches the x-axis at high energies because there is always a small fraction of molecules with very high energies.

The critical feature is the area under the curve to the right of the activation energy, $E_a$. This shaded area represents the fraction of molecules that have enough energy to react when they collide. For the uncatalysed reaction, $E_a$ is high, so only a small fraction of molecules lie to its right. When a catalyst is added, the new lower activation energy $E_{a(cat)}$ sits further to the left on the energy axis — and a much larger area of the curve lies to its right.

The temperature has not changed, so the shape and position of the Maxwell-Boltzmann curve is identical in both cases. The distribution of molecular energies is the same. What has changed is the threshold: the activation energy has moved to a lower value, so a greater proportion of molecules now exceed it. More successful collisions per unit time means a faster reaction rate.

Worked example: Explaining Catalysis Using the Maxwell-Boltzmann Distribution

Explain, with reference to the Maxwell-Boltzmann distribution, how a catalyst increases the rate of a reaction involving a gas. [4 marks]

1. The Maxwell-Boltzmann distribution shows the spread of molecular energies in the gas sample at a given temperature. Most molecules have intermediate energies; very few have very low or very high energies. — [B1] Establishing what the curve represents is the foundation of the explanation. Examiners expect you to name the distribution, not just say 'the graph'.
2. For a reaction to occur, colliding molecules must possess energy greater than or equal to the activation energy, $E_a$. Only the fraction of molecules to the right of $E_a$ on the distribution can react. — [B1] This links the distribution to the concept of activation energy. The phrase 'fraction of molecules' is key exam language.
3. The catalyst provides an alternative reaction pathway with a lower activation energy, $E_{a(cat)}$. The shape of the Maxwell-Boltzmann curve is unchanged because the temperature is the same. — [B1] Stating that the curve is unchanged is important — it shows you understand the catalyst does not heat the system. The curve shifts only with temperature, not with a catalyst.
4. Because $E_{a(cat)} < E_a$, the area under the curve to the right of the activation energy is now larger. A greater proportion of molecules have sufficient energy to react when they collide. — [B1] This is the mark-winning step. Explicitly linking the larger shaded area to a greater proportion of successful collisions is what earns the final mark.
5. More molecules have energy $\geq E_{a(cat)}$ per unit time, so the frequency of successful collisions increases, and the rate of reaction increases. — Closing with 'rate of reaction increases' ties the answer back to the question. Always end with the outcome the question asks about.

Answer: The catalyst provides an alternative pathway with lower activation energy. The Maxwell-Boltzmann distribution is unchanged (same temperature), but the lower $E_a$ means a greater fraction of molecules now exceed the threshold. More successful collisions occur per unit time, so the reaction rate increases.

5. Catalysts in Context: Homogeneous and Heterogeneous

AQA distinguishes between two types of catalyst based on their physical state relative to the reactants. A heterogeneous catalyst is in a different phase from the reactants — for example, solid iron catalysing the gas-phase Haber process. A homogeneous catalyst is in the same phase as the reactants — for example, aqueous Fe$^{2+}$ ions catalysing the reaction between aqueous $\text{S}_2\text{O}_8^{2-}$ and $\text{I}^-$ ions.

In both cases, the fundamental mechanism is the same: the catalyst provides an alternative pathway with lower activation energy. For heterogeneous catalysts, this typically involves adsorption of reactant molecules onto the catalyst surface, weakening bonds and facilitating reaction. For homogeneous catalysts, the catalyst forms an intermediate compound with the reactants, which then breaks down to release the products and regenerate the catalyst.

Regardless of type, the catalyst is regenerated. In the homogeneous Fe$^{2+}$ example, the iron is temporarily oxidised to Fe$^{3+}$ during the reaction, but is reduced back to Fe$^{2+}$ by the end — its chemical composition and amount are unchanged, satisfying the definition precisely.

6. Exam Technique and Mark-Winning Strategies

Questions on catalysts appear across multiple command words in AQA papers. The most demanding are 'explain' questions that require you to link the Maxwell-Boltzmann distribution to the change in activation energy. The step-by-step walkthrough above models the exact structure you need. Here is how to adapt it for different question types.

For a 'describe' question, focus on what happens: the activation energy decreases, the area under the Maxwell-Boltzmann curve to the right of $E_a$ increases, more molecules have sufficient energy to react. Do not include 'because' — that is for 'explain' questions. For an 'evaluate' question about whether a catalyst is beneficial in an industrial process, you must weigh up evidence on both sides and reach a conclusion — for example, considering cost, selectivity, and whether the catalyst is poisoned by impurities.

Catalysts sit at the intersection of kinetics and industrial chemistry throughout the AQA specification. Every time you encounter a rate question — whether it involves temperature, concentration, or surface area — the underlying logic is the same: what fraction of collisions are successful? A catalyst answers that question by lowering the energy threshold, not by changing the energy of the molecules themselves.