Rates of Reaction & Reversible Reactions

What "rate of reaction" means

The rate of reaction tells you how fast reactants turn into products. A fast reaction (like an explosion) finishes in a fraction of a second; a slow one (like iron rusting) can take years.

You can think of rate as the amount of product made (or reactant used up) in a given time:

$$\text{rate} = \frac{\text{amount of product formed}}{\text{time taken}}$$

Key terms Rate of reaction — how quickly reactants are converted into products.

Reactant — a starting substance that is used up during the reaction.

Product — a new substance made by the reaction.

Collision theory

Reactions happen when particles collide. But not every collision causes a reaction. For a collision to be successful (lead to a reaction), two things must be true:

1. The particles must actually collide with each other.
2. They must collide with enough energy to react.

The minimum energy needed for a successful collision is called the activation energy.

Anything that makes particles collide more often or with more energy will increase the rate of reaction. This single idea explains every factor below.

Key terms Collision theory — the idea that particles must collide, with enough energy, for a reaction to occur.

Activation energy — the minimum energy that colliding particles must have for a reaction to happen.

Factors that affect the rate

The table summarises the five factors and how collision theory explains each one.

Exam tip When you explain a rate factor, always mention collisions. A good answer says what changes about the collisions — more frequent collisions, or collisions with more energy, or both. Just saying "particles move faster" without linking it to successful collisions loses marks.

Temperature is special because it does two things at once: it makes collisions more frequent and gives a bigger fraction of particles the activation energy. That is why a small rise in temperature can have a big effect on rate.

Measuring the rate of reaction

To measure a rate you follow something that changes as the reaction goes on. Three common methods:

1. Gas volume given off — collect the gas in a gas syringe (or upturned measuring cylinder over water) and record the volume every few seconds. Good when a gas is produced, e.g. marble chips + acid giving $CO_2$.
2. Loss of mass — stand the flask on a balance; as gas escapes the mass falls. Record the mass at regular times. (A cotton-wool plug stops acid spray escaping but lets gas out.)
3. Cloudiness / disappearing cross — for reactions that form a precipitate, e.g. sodium thiosulfate + hydrochloric acid making cloudy sulfur. Place the flask over a paper cross and time how long until the cross disappears. A shorter time means a faster reaction.

Worked example 0.5 g of marble chips reacts with acid and 24 cm³ of $CO_2$ is collected in 30 s.

rate $= \dfrac{24}{30} = 0.8 \text{ cm}^3/\text{s}$ (average over those 30 s).

Interpreting rate graphs

If you plot product formed (or gas volume) against time, the shape of the curve tells you about the rate:

The steeper the line, the faster the reaction at that moment.
The reaction is fastest at the start (most reactant present, most collisions).
The curve levels off (becomes flat) when a reactant runs out — the reaction has stopped.
A graph that is steeper at the start but levels off at the same height used the same amount of reactant, just faster.

Both curves above finish at the same height because the same quantity of reactant was used — the fast reaction simply gets there sooner and its line is steeper at the start.

Catalysts

A catalyst speeds up a reaction but is not used up itself — you have the same mass of it at the end as at the start, so it can be reused.

A catalyst works by providing a different reaction pathway with a lower activation energy. This means a larger fraction of collisions now have enough energy to succeed, so the rate goes up. The catalyst does not change how much product you make — only how fast you make it.

Real world Catalysts are vital in industry because they cut energy costs and speed up production. Iron is the catalyst in the Haber process, and the catalytic converter in a car uses platinum to speed up the breakdown of harmful exhaust gases.

Reversible reactions

In many reactions the products can react together to re-form the reactants. These are reversible reactions, shown with a double arrow $\rightleftharpoons$ instead of a single arrow.

$$A + B \rightleftharpoons C + D$$

The forward reaction ($A + B \rightarrow C + D$) and the backward reaction ($C + D \rightarrow A + B$) happen at the same time.

Key terms Reversible reaction — one in which the products can react to re-form the reactants; written with the $\rightleftharpoons$ symbol.

Dynamic equilibrium

If a reversible reaction happens in a closed system (nothing can get in or out), it eventually reaches dynamic equilibrium.

At equilibrium:

The forward and backward reactions are still going (it is dynamic, not stopped).
They happen at exactly the same rate.
So the amounts of reactants and products stay constant — they do not change, even though both reactions continue.

Watch out "Equilibrium" does not mean equal amounts of reactants and products, and it does not mean the reactions have stopped. It means the rates of the forward and backward reactions are equal, so concentrations stay constant.

Changing the conditions (qualitative)

Changing the conditions of a system at equilibrium shifts the position of equilibrium — that is, it changes the balance between reactants and products to favour one side.

A useful example is the Haber process, which makes ammonia for fertilisers:

$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$

To get a good yield of ammonia at a sensible rate, industry uses:

a temperature of about 450 °C (a compromise: lower temperature gives more ammonia but too slowly),
a high pressure of about 200 atmospheres (favours the side with fewer gas molecules — the ammonia side),
an iron catalyst (speeds up reaching equilibrium but does not change the position).

Exam tip A catalyst changes how fast equilibrium is reached but never changes the position of equilibrium or the final yield. Only changes to temperature, pressure or concentration shift the position.