What is the rate — and how do we measure it?

The rate of reaction is how quickly the reaction happens. We find it by measuring the loss of a reactant or the gain of a product over time. CCEA gives three practical methods:

• Mass loss — gas escapes, so the flask gets lighter on a balance (e.g. CaCO₃ + acid → CO₂).
• Gas volume — collect the gas in a syringe (e.g. Mg + acid → H₂, or H₂O₂ → O₂).
• Precipitate / turbidity — time how long a mark takes to vanish as the mixture turns cloudy (sodium thiosulfate + acid forms a sulfur precipitate).

Mean (average) rate over a period = quantity changed ÷ time taken — e.g. cm³ of gas per second, or grams lost per second.

Pick the method

Reading a volume–time graph

Plot the volume of gas against time and you get a curve. Its gradient (steepness) is the rate at that moment — the tangent shows it:

Watch out: a graph that levels off higher just made more product — only a steeper curve means a faster rate.

Your turn — mean rate from data

Why reactions happen — collisions

Particles must collide to react. Collision theory says a collision is only successful if the particles meet with enough energy — the activation energy — otherwise they just bounce apart.

So a reaction speeds up if we increase either the frequency of collisions or the energy of collisions:

Watch out: raising the temperature increases both the frequency of collisions and the proportion with enough energy — that’s why heat speeds reactions up so strongly.

What makes a collision count?

The factors that change rate

Each factor works by changing how often, or how energetically, particles collide:

• Temperature ↑ — particles move faster, so collisions are more frequent and more energetic. Rate ↑.
• Concentration ↑ — more particles in the same volume, so collisions are more frequent. Rate ↑.
• Pressure ↑ (gases) — squeezes gas particles closer, like raising concentration. Rate ↑.
• Surface area ↑ (smaller pieces) — a larger surface-area-to-volume ratio exposes more particles to collide. Rate ↑.
• Catalyst — lowers activation energy, so more collisions succeed. Rate ↑.

Does it speed up or slow down?

Tap the option that correctly says what each change does to the rate.

Catalysts — the rate-boosters

A catalyst is a substance that increases the rate of a reaction without being used up — it is chemically unchanged at the end, so a tiny amount works again and again.

It works by providing an alternative reaction pathway with a lower activation energy (the green curve on the earlier diagram). Transition metals and their compounds are often used — e.g. manganese(IV) oxide for H₂O₂, iron in the Haber process.

How a catalyst works

Reversible reactions & dynamic equilibrium

Most GCSE reactions go to completion (one-way arrow). A reversible reaction can go both ways — products can turn back into reactants — shown by the ⇌ sign:

Watch out: equilibrium is dynamic, not stopped. Nothing appears to change, but both reactions carry on at the same rate — and only in a closed system (nothing escapes).

What is dynamic equilibrium?

Le Châtelier’s principle

Le Châtelier’s principle: if a change is made to a system at equilibrium, the position of equilibrium shifts to oppose the change.

For a closed, homogeneous system, the qualitative effects are:

• Concentration: add more of a substance → equilibrium shifts away from it (to use it up). Remove it → shifts towards it (to replace it).
• Temperature: heat it → shifts in the endothermic direction (to absorb the heat). Cool it → shifts in the exothermic direction.
• Pressure (gases): raise pressure → shifts to the side with fewer moles of gas (to reduce pressure). Lower pressure → shifts to the side with more moles of gas.

Watch out: equilibrium always shifts to oppose whatever you did — add heat, it absorbs heat; squeeze it, it makes fewer gas molecules. A catalyst speeds both rates equally, so it does not move the position of equilibrium.

Predict the shift

For A + B ⇌ C + D (forward = exothermic), tap a change on the left, then its correct shift on the right.

Apply Le Châtelier

The Haber process — a real trade-off

Ammonia is made from nitrogen and hydrogen in a reversible reaction:

Le Châtelier says a high pressure (4 gas moles → 2) and a low temperature (exothermic) give the best yield. But a low temperature gives a very slow rate, so industry compromises:

• Temperature ≈ 450 °C — a compromise: lower yield, but a workable rate.
• Pressure ≈ 200 atmospheres — high yield, but kept safe and affordable.
• Iron catalyst — speeds up the rate (reaches equilibrium faster) without shifting its position.

Key idea: the conditions chosen balance a good rate of production against a good position of equilibrium (yield) at a reasonable cost.

Choosing Haber conditions

The points to know

You’ve covered all of CCEA Units 2.3 & 2.4 — rates, collision theory, catalysts, reversible reactions, equilibrium, Le Châtelier and the Haber process. Press Finish to see your score.