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This lesson covers reversible reactions, dynamic equilibrium and Le Chatelier's principle — all of which are Higher tier topics in the Edexcel GCSE Chemistry (1CH0) specification. Understanding these concepts is essential for explaining how industrial processes such as the Haber process are optimised. This is one of the most challenging topics in GCSE Chemistry but also one of the most rewarding once mastered.
Note: This topic is assessed at Higher tier only. Foundation tier students are not required to understand dynamic equilibrium or Le Chatelier's principle, but they should know that some reactions are reversible.
A reversible reaction is a reaction that can proceed in both the forward direction (reactants → products) and the backward (reverse) direction (products → reactants).
The symbol for a reversible reaction is ⇌ (two half-arrows pointing in opposite directions).
The reaction between nitrogen and hydrogen to form ammonia:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Both reactions occur simultaneously. The system never completely converts to products — there is always a mixture of reactants and products present.
CuSO₄(s) + 5H₂O(l) ⇌ CuSO₄·5H₂O(s)
Exam tip: The ⇌ symbol is very important. If you see it in an equation, it tells you the reaction is reversible. Never use a single arrow (→) for a reversible reaction.
When a reversible reaction takes place in a closed system (one where no substances can enter or leave), a state of dynamic equilibrium is eventually reached.
At dynamic equilibrium:
| Term | Meaning |
|---|---|
| Dynamic | Both forward and reverse reactions are still happening |
| Equilibrium | The rates of the forward and reverse reactions are equal |
| Closed system | No substances can enter or leave the reaction vessel |
| Constant concentrations | The amounts of reactants and products do not change over time (but they are not necessarily equal) |
Exam tip: Many students mistakenly think that equilibrium means "the reaction has stopped." This is wrong. At dynamic equilibrium, both reactions are still occurring — they are just happening at the same rate, so there is no overall change in concentrations.
Le Chatelier's principle states:
If a system at equilibrium is subjected to a change in conditions, the equilibrium will shift in the direction that tends to counteract (oppose) the change.
In other words, if you disturb an equilibrium, the system adjusts to partially undo what you did.
This principle allows us to predict how changes in temperature, concentration and pressure affect the position of equilibrium.
For the general reversible reaction:
A + B ⇌ C + D
If the forward reaction is exothermic (releases heat), then the reverse reaction is endothermic (absorbs heat).
| Change in temperature | Direction of shift | Explanation |
|---|---|---|
| Increase temperature | Shifts towards the endothermic direction | The system absorbs the extra heat to counteract the increase |
| Decrease temperature | Shifts towards the exothermic direction | The system releases heat to counteract the decrease |
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) forward reaction is exothermic (ΔH = −92 kJ/mol)
However, if the temperature is too low, the rate of reaction becomes very slow, even though the equilibrium position favours products. This is why a compromise temperature is chosen.
Exam tip: Changing the temperature changes both the position of equilibrium and the rate of reaction. A higher temperature always increases the rate (particles have more energy), but it may shift the equilibrium away from the desired product. This is why a compromise is needed.
| Change in concentration | Direction of shift | Explanation |
|---|---|---|
| Increase concentration of a reactant | Shifts to the right (towards products) | The system uses up the extra reactant |
| Decrease concentration of a reactant | Shifts to the left (towards reactants) | The system replaces the lost reactant |
| Increase concentration of a product | Shifts to the left (towards reactants) | The system uses up the extra product |
| Decrease concentration of a product | Shifts to the right (towards products) | The system replaces the lost product |
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Pressure only affects equilibria involving gases. Le Chatelier's principle tells us:
| Change in pressure | Direction of shift | Explanation |
|---|---|---|
| Increase pressure | Shifts towards the side with fewer gas molecules | Fewer molecules exert less pressure, counteracting the increase |
| Decrease pressure | Shifts towards the side with more gas molecules | More molecules exert more pressure, counteracting the decrease |
To apply this, count the number of moles of gas on each side of the equation.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
| Change | Shift | Effect on NH₃ yield |
|---|---|---|
| Increase pressure | Towards the right (fewer gas molecules) | More NH₃ at equilibrium |
| Decrease pressure | Towards the left (more gas molecules) | Less NH₃ at equilibrium |
So high pressure favours the production of ammonia. However, very high pressures are expensive and dangerous (strong equipment is needed), so a compromise pressure of about 200 atmospheres is used.
A catalyst does NOT change the position of equilibrium.
A catalyst:
Exam tip: This is a very commonly tested point. Catalysts speed up the attainment of equilibrium but do NOT shift the equilibrium position. If asked "does a catalyst change the yield of product?" the answer is no — it only changes how quickly equilibrium is reached.
The Haber process is used to manufacture ammonia (NH₃) from nitrogen and hydrogen:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol (exothermic forward reaction)
The raw materials are nitrogen (from fractional distillation of liquid air) and hydrogen (from the reaction of natural gas with steam).
The conditions used in the Haber process represent a compromise between yield and rate:
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