3cReversible reactions and equilibria

Most reactions run one way and stop. Some can run both ways, and those are written with a double arrow:

A + B ⇌ C + D

The ⇌ symbol is the signal. A single arrow means the reaction goes one way only.

What reversible means in practice

The products can react together to re-form the reactants. So in a closed container a reversible reaction never goes to completion — some of every substance is always present.

Hydrated copper(II) sulfate

The standard example, and one you can watch both ways.

Heating blue hydrated copper(II) sulfate drives off the water of crystallisation, leaving white anhydrous copper(II) sulfate.

Adding water to the white solid turns it blue again, and releases heat.

That reverse reaction is exothermic — the mixture becomes noticeably warm, and it is a creditable observation.

The sharp colour change in both directions is what makes anhydrous copper(II) sulfate the standard test for water.

Ammonium chloride

Heat solid ammonium chloride and it decomposes into two gases: ammonia and hydrogen chloride.

Where the tube is cooler, those gases recombine and a white solid re-forms on the glass.

This is not sublimation. The solid genuinely decomposes and then re-forms — a chemical change happening twice, in opposite directions.

Why an open container changes things

Heat hydrated copper(II) sulfate in an open test tube and the water vapour escapes. With no water left, the reverse reaction cannot happen.

So a reversible reaction in an open container behaves as though it goes to completion. Losing a product prevents the reverse reaction entirely — which is exactly why equilibrium needs a sealed container.

Separate Chemistry only.

In a sealed container, a reversible reaction settles into a state where nothing appears to change. That state is dynamic equilibrium.

The two characteristics

1. The forward and reverse reactions happen at the same rate.
2. The concentrations of reactants and products stay constant.

Both are needed for a full definition. Giving only the second misses the word dynamic.

Nothing has stopped

The commonest misconception is that equilibrium means the reaction has finished.

It has not. Both reactions continue, at rates that exactly cancel. Products are being formed as fast as they are converted back, so the concentrations hold steady while the chemistry carries on underneath.

That is what dynamic means, and it is why the word is in the name.

Constant is not equal

At equilibrium the concentrations stop changing. They are almost never equal to each other.

A mixture can sit at equilibrium with far more product than reactant — that simply means the position of equilibrium lies towards the products. Equal rates are required; equal amounts are not.

Why the container must be sealed

If a product escapes, the reverse reaction is starved and can never match the forward one. The concentrations never become constant, so equilibrium is never reached.

This is why a reversible reaction producing a gas, carried out in an open beaker, simply runs to completion.

Catalysts and equilibrium

A catalyst speeds up the forward reaction and the reverse reaction by the same amount.

So the two rates stay equal to each other, and the position of equilibrium does not move. Equilibrium is simply reached sooner.

That distinction — faster to arrive, same destination — is what these questions test. In 4CH1 the factors you are asked to apply to the position of equilibrium are temperature and pressure.

Concentration also shifts an equilibrium, and you may meet that in later study — but the questions here are built around temperature and pressure, and a catalyst is the one that changes only the speed.

Separate Chemistry only.

Two conditions move the position of equilibrium, and each is decided by a separate rule.

Temperature

Increasing the temperature favours the endothermic direction. Decreasing it favours the exothermic direction.

So work out which direction is endothermic first, then decide which way the equilibrium shifts.

For N₂ + 3H₂ ⇌ 2NH₃, the forward reaction is exothermic. So lowering the temperature shifts the equilibrium towards ammonia and raises the yield.

Pressure

Increasing the pressure favours the side with fewer gas molecules.

Count the molecules on each side, including the coefficients.

For N₂ + 3H₂ ⇌ 2NH₃: that is 1 + 3 = 4 on the left, 2 on the right. Higher pressure shifts it right, raising the ammonia yield.

If both sides have the same number of gas molecules — as in H₂ + I₂ ⇌ 2HI, two and two — pressure has no effect on the position at all.

Always count first.

The two rules are independent

Temperature is decided by which direction is endothermic. Pressure is decided by which side has fewer gas molecules. They can point the same way or opposite ways, and each is worked out separately.

Why industry compromises

Ammonia is made at about 450 °C, even though a lower temperature would give a higher yield.

The reason is that the two things you want pull in opposite directions:

• A low temperature gives a better yield.
• A high temperature gives a faster rate.

At a low temperature the reaction would be too slow to be economic. 450 °C is a compromise — an acceptable yield in an acceptable time.

The same logic applies to pressure. Higher pressure would raise the yield further, but stronger vessels cost more, compressing gases uses energy, and very high pressures are dangerous. About 200 atmospheres balances chemistry against cost and safety.

When a question asks why particular industrial conditions are used, the word it wants is compromise — and a full answer says what is gained and what is given up.