Many chemical reactions run in both directions: products can re-form the original reactants. When forward and reverse rates become equal in a closed system, the mixture reaches dynamic equilibrium — concentrations stop changing even though both reactions continue. Le Chatelier's principle predicts how the system responds when you disturb that balance.
What is a reversible reaction?
A reversible reaction is one in which the products can react to re-form the original reactants. The double arrow (⇌) indicates a reversible reaction:
A + B ⇌ C + D
A classic classroom example is the reaction of hydrated copper(II) sulfate:
CuSO₄·5H₂O (blue solid) ⇌ CuSO₄ (white solid) + 5H₂O (steam)
Heating drives the reaction to the right (forward reaction); adding water drives it to the left (reverse reaction).
What is dynamic equilibrium?
Dynamic equilibrium is reached in a closed system (no substances can enter or leave) when the rate of the forward reaction equals the rate of the reverse reaction. At this point:
- The concentrations of reactants and products are constant (not necessarily equal).
- Both the forward and reverse reactions are still occurring — the equilibrium is dynamic (not static).
- The system looks unchanged macroscopically, but at the molecular level both reactions are still happening.
The position of equilibrium describes whether there is more product or more reactant at equilibrium:
- Equilibrium to the right → more products than reactants (favours forward reaction)
- Equilibrium to the left → more reactants than products (favours reverse reaction)
What is Le Chatelier's principle?
Le Chatelier's principle (1884) states:
"If a system at equilibrium is subjected to a change in conditions, the equilibrium shifts in the direction that opposes the change."
This principle allows you to predict the effect of changing three main variables: concentration, temperature, and pressure (for gas reactions).
How do changes in concentration affect equilibrium?
| Change | Equilibrium shift | Effect |
|---|---|---|
| Increase concentration of reactant | Shifts right (forward) | More product formed |
| Decrease concentration of reactant | Shifts left (reverse) | More reactant formed |
| Increase concentration of product | Shifts left (reverse) | Products are consumed |
| Decrease concentration of product | Shifts right (forward) | More product formed |
In industrial processes, removing products as they form (e.g. by condensation or absorption) continuously shifts the equilibrium to the right, increasing overall yield.
How does temperature affect equilibrium?
Changing temperature shifts the equilibrium in the direction of the endothermic reaction:
| Change | Equilibrium shift | Effect on yield of products |
|---|---|---|
| Increase temperature | Shifts toward endothermic direction | Favours the reaction that absorbs heat |
| Decrease temperature | Shifts toward exothermic direction | Favours the reaction that releases heat |
Note: changing temperature changes the actual equilibrium constant (K) — the ratio of product to reactant concentrations at equilibrium shifts to a new value. This is different from changing concentration or pressure, which shift the position but not the value of K.
How does pressure affect equilibrium in gas reactions?
Changing pressure only affects equilibria that involve gases. The equilibrium shifts toward the side with fewer moles of gas:
| Change | Equilibrium shift |
|---|---|
| Increase pressure | Shifts toward side with fewer gas molecules |
| Decrease pressure | Shifts toward side with more gas molecules |
Worked example — Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Count moles of gas: left side has 1 + 3 = 4 moles; right side has 2 moles.
- Increasing pressure → shifts right → more ammonia produced ✓
- Decreasing pressure → shifts left → less ammonia produced ✗
High pressure therefore favours ammonia production.
What is the Haber process and why are the conditions a compromise?
The Haber process manufactures ammonia for fertilisers and explosives:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹ (exothermic)
Applying Le Chatelier's principle:
| Variable | What maximises yield | Practical problem |
|---|---|---|
| Temperature | Low temperature (forward reaction is exothermic, so cooling shifts right) | At low temperature the reaction is too slow — not economically viable |
| Pressure | High pressure (fewer moles of gas on product side) | Very high pressures are expensive and dangerous |
Industrial compromise:
- Temperature: ~450 °C (fast enough rate; acceptable — though not maximum — yield)
- Pressure: ~200 atmospheres (high enough for good yield; economically feasible)
- Catalyst: iron + potassium hydroxide promoter (increases rate without affecting position of equilibrium)
- Continuous removal of ammonia by condensation (shifts equilibrium right; unreacted N₂ and H₂ are recycled)
The yield at these conditions is only about 15 %, but continuous recycling makes the overall process efficient.
Frequently asked questions
What is the difference between a reversible reaction and dynamic equilibrium?
A reversible reaction is any reaction that can proceed in both directions. Dynamic equilibrium is a specific state reached in a closed system when the rate of the forward reaction equals the rate of the reverse reaction, so concentrations no longer change. All systems at dynamic equilibrium involve reversible reactions, but not all reversible reactions reach equilibrium — if the container is open and products escape, equilibrium is never established.
Why does a catalyst not change the position of equilibrium?
A catalyst speeds up both the forward and reverse reactions equally, so it does not change the concentrations of reactants or products at equilibrium — the position of equilibrium is unchanged. What the catalyst does is allow the system to reach equilibrium faster. In industrial applications like the Haber process, the catalyst is economically vital because it allows an acceptable rate of production at a lower (and more practical) temperature.
Why is high temperature used in the Haber process if low temperature gives a better yield?
At low temperature the reaction is so slow that the rate of ammonia production would be uneconomically low — the nitrogen and hydrogen molecules do not have enough energy to react frequently enough. The industrial process uses 450 °C as a compromise: this gives a sufficiently fast reaction rate to produce ammonia quickly, while still giving a reasonable (if not maximum) equilibrium yield. The use of an iron catalyst also helps achieve an acceptable rate at this intermediate temperature.
What does it mean if we say equilibrium lies "to the right"?
When we say equilibrium lies to the right, we mean that at equilibrium there is a greater concentration of products than reactants. A reaction equation is conventionally written with reactants on the left and products on the right, so "lying to the right" means the forward reaction is favoured — more products are present at the equilibrium point than reactants. If equilibrium lies to the left, reactants predominate and the forward reaction is not very efficient.
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