What "rate of reaction" actually means

The rate of reaction is the change in the concentration (or amount) of a reactant or product per unit time.

Anything that changes measurably as the reaction proceeds can be used to follow it:

• Gas volume collected in a gas syringe (e.g. H₂ from Mg + HCl).
• Mass loss on a balance as a gas escapes (e.g. CO₂ from CaCO₃ + HCl).
• Colour change followed with a colorimeter (e.g. the fading purple of MnO₄⁻).
• Titrating samples quenched at known times, or measuring pH / conductivity.

Mean rate vs initial rate. A mean rate divides the total change by the total time. The initial rate is the gradient of the tangent at t = 0 — it is the fastest rate, because the concentrations are highest then and nothing has been used up.

Reading a rate off a graph

Your turn — mean rate from gas volume

Your turn — mean rate from mass loss

Your turn — rate in mol dm⁻³ s⁻¹

Your turn — initial rate from a tangent

Turning moles into a rate

Exam questions often make you do the mole arithmetic before you can find the rate. Remember: at room temperature and pressure, 1 mol of any gas occupies 24 000 cm³ (24.0 dm³).

Your turn — volume of hydrogen

Your turn — and the rate

Collision theory and activation energy

For particles to react they must collide — but the overwhelming majority of collisions achieve nothing. A collision is successful only if:

• the colliding particles have a combined energy greater than or equal to the activation energy, E_a; and
• they collide in the correct orientation (the reacting parts of the molecules must meet).

So the rate depends on the frequency of successful collisions — and every single factor below works by changing either how often particles collide, or what fraction of those collisions has E ≥ E_a.

A high E_a means a slow reaction at a given temperature, because only a tiny fraction of molecules have enough energy. This is why a mixture of petrol vapour and air sits happily in a fuel tank until a spark supplies the activation energy.

Defining activation energy

The Maxwell–Boltzmann distribution

At any temperature the molecules in a gas do not all have the same energy. The Maxwell–Boltzmann distribution shows how molecular energies are spread out.

• The peak is the most probable energy — not the mean. The mean energy is slightly to the right of the peak because the curve is not symmetrical.
• The total area under the curve = the total number of molecules.
• The area beyond E_a = the number of molecules with enough energy to react. It is a tiny fraction of the total.

What temperature does to the distribution

Raise the temperature and the whole distribution changes shape:

• The molecules have more kinetic energy, so they move faster and collide more frequently.
• Far more importantly, a much greater proportion of molecules now has E ≥ E_a — the area beyond E_a grows dramatically.

Exam gold: a 10 K rise typically doubles the rate, yet it raises the mean speed by only about 2%. That is far too small to explain a doubling by "more frequent collisions" alone — the increase comes overwhelmingly from the fraction of molecules exceeding E_a. Say this and you get the mark.

Why heating works

Reading the curve

Concentration, pressure and surface area

• Concentration ↑ — more particles in the same volume, so the particles are closer together and collide more frequently. The proportion of successful collisions is unchanged; there are simply more collisions per second.
• Pressure ↑ (gases, at constant temperature) — squeezing the same number of molecules into a smaller volume raises the concentration of the gases. Same explanation: more frequent collisions.
• Surface area ↑ (a solid reactant) — grinding a lump into a powder exposes more particles at the surface where the other reactant can reach them, so collisions are more frequent. The mass is the same; the accessible surface is not.

Two traps. (1) Adding an inert gas at constant volume raises the total pressure but does not change the concentration of the reactants — so the rate is unchanged. (2) Pressure has essentially no effect on reactions between liquids or solutions, which are almost incompressible.

An inert gas

Catalysts and activation energy

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy. It is not used up — it is regenerated at the end, so a tiny amount catalyses a huge quantity of reaction.

• A lower E_a means a much larger proportion of molecules has enough energy to react — draw a second dashed line further left on your Maxwell–Boltzmann curve and the shaded area leaps up.
• Heterogeneous catalyst — in a different phase from the reactants (e.g. solid Fe in the Haber process); reaction happens on the surface.
• Homogeneous catalyst — in the same phase (e.g. Cl• radicals catalysing ozone breakdown in the gas phase).

What a catalyst does NOT do: it does not change ΔH, it does not shift the position of an equilibrium, and it does not "lower the energy of the reactants". It offers a different route.

What a catalyst does

Why powders react faster

Faster, slower — or no change?

A reaction between two gases is running in a sealed, rigid vessel, alongside a solid + acid reaction on the bench. Tap a change, then tap what it does to the rate.

Match the factor to the explanation

Tap an item on the left, then its partner on the right.

The big ideas to know

You've now covered Topic 9: Kinetics I of the Edexcel A-level Chemistry (9CH0) specification. Press Finish to see your score.