Reversible Reactions
So far every reaction has gone one way: reactants → products, done. But in some reactions the products can react together to re-form the original reactants. These are reversible reactions, and they get their own arrow:
A + B ⇌ C + D
The double half-arrow ⇌ means the reaction goes in both directions at the same time: A and B react to make C and D (the forward reaction), while C and D react to re-form A and B (the reverse reaction). Which direction is favoured can be controlled by changing the conditions — temperature, pressure or concentration.
A classic example is heating ammonium chloride. Warm it and the white solid breaks down into two gases; let the gases cool and they recombine into the white solid — which is why a heated tube of ammonium chloride grows a white ring further up, where the gases meet cooler glass:
ammonium chloride ⇌ ammonia + hydrogen chloride
NH4Cl(s) ⇌ NH3(g) + HCl(g)
White ammonium chloride solid sits in the bottom of the tube.
Heating pushes the reaction forward (decomposition) at the hot bottom of the tube; higher up, the cooler glass pushes it back (recombination) — the same reversible reaction running both ways at once.
Energy changes in reversible reactions
If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction — and the same amount of energy is transferred in each case. Energy is conserved: whatever the forward reaction gives out, the reverse reaction must take back in.
Some questions give the energy change as ΔH with a sign instead of the words “exothermic” or “endothermic”, and you need to read which one it means. A negative ΔH (e.g. ΔH = −45 kJ/mol) means the forward reaction is exothermic; a positive ΔH means it is endothermic. The reverse reaction is then the opposite, transferring the same amount of energy. You won’t be asked to work out a ΔH value at GCSE — just to read what its sign tells you.
You can see this in the ammonium chloride reaction above: the forward direction (decomposition) only happens while you keep heating it, so it is endothermic — which means the reverse change (the gases recombining into the white ring on the cooler glass) must be exothermic, releasing exactly the same amount of energy.
🧪 Exam-style questions
What does the symbol ⇌ in an equation tell you? Tick (✓) one box.
The forward direction of a reversible reaction is exothermic and transfers 58 kJ of energy. How much energy is transferred by the reverse reaction?
Show answer
- 58 kJ — taken in (endothermic). 1 mark
- If a reversible reaction is exothermic one way, it is endothermic the other way, and the same amount of energy is transferred in each case.
For a reversible reaction, the forward direction has ΔH = −45 kJ/mol. Which statement is correct? Tick (✓) one box.
Heating ammonium chloride in a test tube produces a white ring of ammonium chloride higher up the tube. Explain why.
Show a model answer
- Heating decomposes the ammonium chloride into ammonia and hydrogen chloride gases (the forward, endothermic direction). 1 mark
- The gases rise to a cooler part of the tube, where the reverse reaction happens: they recombine into solid ammonium chloride — the white ring. 1 mark
- One experiment, both directions: changing the conditions (here, temperature) changes which way a reversible reaction is favoured.
Dynamic Equilibrium
Run a reversible reaction in an open beaker and the gaseous products drift away — the reverse reaction never gets its ingredients. But seal the same reaction into a closed system (“apparatus which prevents the escape of reactants and products”) and something remarkable happens.
At first the forward reaction is fast — the flask is full of reactants — and the reverse reaction can barely happen, because there are hardly any products to react. As reactants are used up, the forward rate falls; as products accumulate, the reverse rate rises. Eventually the two rates meet:
When a reversible reaction occurs in apparatus which prevents the escape of reactants and products, equilibrium is reached when the forward and reverse reactions occur at exactly the same rate.
Approaching equilibrium in a closed system: the forward rate falls as reactants are used up, the reverse rate climbs as products build, and at equilibrium the two are exactly equal — both reactions still running, cancelling each other out.
The same approach to equilibrium looks just as clear if we plot the concentrations against time instead of the rates:
Now plotted as concentration against time. The reactants fall and the products rise until, at equilibrium, every concentration goes constant — the lines flatten together. But notice they settle at different heights: at equilibrium the amounts are constant, not equal. How high each one settles depends on where the position of equilibrium lies.
That is why it is called dynamic equilibrium: nothing has stopped. Both reactions are still running, but because they run at the same rate, the concentrations of all the substances stay constant. From the outside the mixture looks frozen; at the particle level it is anything but.
- “At equilibrium the reaction stops.” No — both reactions continue. The rates are equal, so there is no further overall change. Concentrations are constant, not static because nothing is happening.
- “At equilibrium there are equal amounts of reactants and products.” No — the amounts are constant, not equal. An equilibrium can sit far to the right (mostly products) or far to the left (mostly reactants). Where it sits is called the position of equilibrium — and moving that position is the whole of section 7.
Equilibrium can only be reached in a closed system. If a product escapes — a gas bubbling out of an open flask — the reverse reaction is starved of its reactant, the reverse rate can never catch up with the forward rate, and the reaction simply runs to completion. (The heated ammonium chloride tube in section 5 only forms its white ring because the tube traps the gases long enough for them to recombine.)
Some equilibria involve a coloured substance, and the colour reveals where the position lies. A classic is hydrogen + iodine ⇌ hydrogen iodide, H2(g) + I2(g) ⇌ 2HI(g): iodine vapour is purple, while H2 and HI are colourless. As the forward reaction proceeds the mixture fades to a paler purple — but it never goes colourless, because at equilibrium some I2 always remains (the reaction doesn’t go to completion). “Paler, not colourless” is the mark-scheme distinction.
The colour reveals the position
Drag the slider to move the position of equilibrium left or right, and watch the colour change as the amounts change.
H2(g) + I2(g) ⇌ 2HI(g)
I2 vapour is purple; H2 and HI are colourless — so the colour you see tracks the amount of I2. However far right the position lies, a little I2 always remains: paler, never colourless.
🧪 Exam-style questions
Which statement is true for a reversible reaction at equilibrium? Tick (✓) one box.
Why can equilibrium only be reached in a closed system? Tick (✓) one box.
At equilibrium, the concentrations of all reactants and products are… Tick (✓) one box.
Colourless hydrogen and purple iodine vapour are sealed in a tube and left to reach equilibrium: H2(g) + I2(g) ⇌ 2HI(g). What is the colour of the mixture at equilibrium? Tick (✓) one box.