Collision theory and activation energy, the effects of concentration, pressure, surface area, temperature and catalysts on rate, the Maxwell-Boltzmann distribution, and how a catalyst provides an alternative pathway of lower activation energy.

What this dot point is asking

CCEA wants you to use collision theory to explain reaction rate, define activation energy, explain the effects of concentration, pressure, surface area, temperature and catalysts, sketch and interpret the Maxwell-Boltzmann distribution, and explain how a catalyst increases rate by lowering activation energy.

Collision theory and activation energy

The rate depends on the frequency of these successful collisions.

Factors affecting rate

• Concentration and pressure: more particles per unit volume, so more frequent collisions and a higher rate.
• Surface area: breaking a solid into smaller pieces exposes more particles, raising collision frequency.
• Temperature: particles move faster (more frequent collisions) and a greater proportion exceed $E_a$, which is the dominant effect.

For concentration, pressure and surface area, only the collision frequency changes; the proportion of collisions that succeed stays the same because the energy distribution of the molecules is unchanged. Temperature is different and far more powerful, because it changes both the collision frequency and, much more importantly, the proportion of molecules with energy at or above $E_a$. This is why a rough rule of thumb is that a $10\ \text{C}$ rise in temperature roughly doubles the rate, whereas doubling the concentration only doubles the collision frequency. Measuring a rate experimentally usually means following a quantity that changes with time, such as the volume of gas produced, the mass lost, or the time taken for a fixed amount of product (the clock reaction), then comparing initial gradients.

The Maxwell-Boltzmann distribution

Catalysts

On a Maxwell-Boltzmann diagram, the catalyst moves $E_a$ to a lower value, so a larger proportion of molecules now have enough energy to react. The catalyst does not change the enthalpy change. Catalysts may be homogeneous (in the same phase as the reactants, like the chlorine radicals that catalyse ozone breakdown) or heterogeneous (in a different phase, like the solid iron in the Haber process, on whose surface the reaction takes place). In both cases the principle is the same: a new route with a lower activation energy.

Examples in context

Example 1. Catalytic converters in cars. A catalytic converter uses a heterogeneous catalyst (platinum, palladium and rhodium spread over a honeycomb to maximise surface area) to speed up reactions that turn toxic carbon monoxide and nitrogen oxides into carbon dioxide and nitrogen. The catalyst lowers the activation energy of these reactions so they proceed fast enough in the few seconds the gases spend in the exhaust. The large surface area is a deliberate design choice, the same principle as powdering a solid reactant.

Example 2. Enzymes as biological catalysts. Enzymes lower the activation energy of reactions in living cells, allowing them to proceed quickly at body temperature (about $37\ \text{C}$) instead of needing high heat. Catalase, for example, speeds the breakdown of hydrogen peroxide into water and oxygen by an enormous factor. This is a clear demonstration that a catalyst lets a reaction reach a workable rate at a low temperature, exactly the trade-off CCEA expects candidates to articulate.

Try this

Q1. State the two conditions for a successful collision. [2 marks]

• Cue. Energy at or above $E_a$, and the correct orientation.

Q2. Explain how a catalyst increases the rate of reaction. [2 marks]

• Cue. It provides an alternative pathway with a lower $E_a$, so a greater proportion of molecules can react.