Three ways to measure a rate

The rate of reaction tells you how quickly reactants are used up or products are made. To measure it you follow a quantity that changes over time:

• Volume of gas given off — collected in a gas syringe (units cm³).
• Mass lost — a flask on a balance loses mass as gas escapes (units g).
• Precipitation / turbidity — a solution turns cloudy; time how long until a cross under the beaker disappears.

Pick the method

Mean rate of reaction

Over a time interval, the mean rate is simply the quantity made (or used up) divided by the time taken:

Units depend on what you measured: cm³/s for gas volume, g/s for mass loss, or mol/s for moles.

Your turn — mean rate

Reading a rate graph

Plot the quantity (e.g. gas volume) against time and the curve tells the story. The reaction is fastest at the start (steepest gradient), then slows as reactants are used up, and levels off (flat) when the reaction is finished.

Watch out: a steeper curve means a faster rate — not a "bigger" reaction. Two curves can level off at the same height (same amount of product) but one gets there faster.

Rate from a tangent

Why reactions happen

Collision theory says a reaction only happens when particles collide with each other and with enough energy. That minimum energy needed for a collision to be successful is the activation energy (E_a).

Anything that increases the frequency of collisions, or the energy of collisions, increases the rate of reaction.

What is activation energy?

Five ways to speed it up

AQA expects you to explain each factor using collision theory — what it does to the frequency and/or energy of collisions:

• Concentration (of solutions) — more particles in the same volume → more frequent collisions.
• Pressure (of gases) — squeezes particles closer → more frequent collisions.
• Surface area (of solids) — smaller pieces (e.g. powder) expose more surface → more frequent collisions.
• Temperature — particles move faster, so collisions are more frequent and have more energy.
• A catalyst — provides a lower-activation-energy pathway, so more collisions succeed.

Watch out: raising the temperature increases both the frequency and the energy of collisions — many students only mention one. The energy effect is the bigger reason temperature speeds reactions up so much.

Does it increase the rate?

Tap whether each change would increase or not increase the rate of a reaction.

Catalysts & enzymes

A catalyst speeds up a reaction by providing a different pathway with a lower activation energy — but it is not used up in the reaction, so it does not appear in the chemical equation.

• Different reactions need different catalysts.
• Enzymes are biological catalysts — they catalyse the reactions in living things.
• A catalyst lets you use a lower temperature (saving energy and money) to get the same rate.

Watch out: a catalyst lowers the activation energy and is not used up — it does not "give the particles more energy" and it is not a reactant.

How a catalyst works

Investigating rate

The AQA required practical investigates how a factor changes the rate of reaction. You can vary either:

• Concentration — e.g. sodium thiosulfate + hydrochloric acid: change the acid concentration and time the disappearing cross (turbidity).
• Surface area — e.g. marble chips + hydrochloric acid: compare large chips with powder, measuring gas volume or mass loss.

Keep all other variables the same (control variables) so the change you see is caused only by the factor you are testing. A faster rate = shorter time (or steeper graph).

Exam tip: control variables here include temperature, total volume, and the amount of solid — only the tested factor should change.

Fair test

Reactions that go both ways

In a reversible reaction the products can react to re-form the reactants. We show this with the ⇌ symbol:

If a reversible reaction is exothermic in one direction, it is endothermic in the other, by the same amount of energy.

Energy both ways

Equilibrium in a closed system

In a closed system (nothing in or out), a reversible reaction reaches dynamic equilibrium: the forward and backward reactions happen at exactly the same rate. The amounts of reactants and products then stay constant.

Watch out: at equilibrium the reaction has not stopped — the forward and backward reactions continue at equal rates. It is dynamic, not static.

What is happening at equilibrium?

Shifting the equilibrium

Higher tier only. Le Chatelier's principle: if you change the conditions of a system at equilibrium, the position of equilibrium shifts to oppose that change.

• Concentration — increase a reactant: equilibrium shifts right (towards products) to use it up.
• Temperature — increase it: equilibrium shifts in the endothermic direction (to absorb the added energy).
• Pressure (gases) — increase it: equilibrium shifts to the side with fewer molecules of gas.

Watch out: the system opposes the change you make. Raise the temperature and it favours the endothermic direction; add more of a substance and it shifts away from that substance.

Predict the shift

Change & response

Match each change to the way the equilibrium responds. Tap one on the left, then its partner on the right.

The key ideas to know

You've covered the whole of AQA 4.6 — rate of reaction, collision theory, factors, catalysts, the required practical, reversible reactions and equilibrium. Press Finish to see your score.