Chemical equilibria
Section: 7 Equilibria | Syllabus: Cambridge AS Level Chemistry (9701)
Reversible Reactions
Reversible Reaction A reaction in which the products can themselves react to reform the original reactants. Represented using the reversible arrow symbol (⇌).In a reversible reaction, two processes occur simultaneously:
- Forward reaction: reactants → products
- Reverse reaction: products → reactants
NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
Heating decomposes it into gases; on cooling the gases recombine to reform the solid.
Other common reversible reactions:
- CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(g) - blue hydrated ⇌ white anhydrous copper(II) sulfate
- N₂(g) + 3H₂(g) ⇌ 2NH₃(g) - the Haber process
- 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) - the Contact process
Dynamic Equilibrium
Dynamic Equilibrium The state reached in a closed system when the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant over time.Two essential features of dynamic equilibrium:
- Constant concentrations: the amounts of reactants and products do not change over time
- Dynamic - not static: both forward and reverse reactions continue to occur; they simply proceed at the same rate
Diagram: Graph of reaction rate on the y-axis against time on the x-axis. The forward reaction rate starts high and decreases; the reverse reaction rate starts at zero and increases. The two lines meet and become equal at the point of dynamic equilibrium, marked with a vertical dashed line. Both rates remain constant after that point.
Closed System A system in which no matter can enter or leave, although energy can be exchanged with the surroundings. Dynamic equilibrium can only be established in a closed system - if the system is open, products or reactants escape before equilibrium is reached. Common Misconception Equilibrium does not mean equal concentrations of reactants and products. It means concentrations are constant (no longer changing). The position of equilibrium - whether it favours reactants or products - depends on the specific reaction.Le Chatelier's Principle
Le Chatelier's Principle If a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise that change.The "position of equilibrium" describes the relative amounts of reactants and products in the equilibrium mixture:
- Shifts right (→): more products form (forward reaction favoured)
- Shifts left (←): more reactants form (reverse reaction favoured)
Effect of Changing Concentration
| Change | Shift Direction | Reason |
|---|---|---|
| Increase [reactant] | Right (→) | System consumes the added reactant by producing more product |
| Decrease [reactant] | Left (←) | System replaces lost reactant by decomposing product |
| Increase [product] | Left (←) | System removes the added product by reforming reactant |
| Decrease [product] | Right (→) | System replaces lost product by producing more |
Effect of Changing Temperature
For an equilibrium where the forward reaction is exothermic (releases heat), the reverse reaction is endothermic (absorbs heat).
| Change | Shift Direction | Effect on K |
|---|---|---|
| Increase temperature | Endothermic direction - absorbs heat to minimise the increase | K changes |
| Decrease temperature | Exothermic direction - releases heat to minimise the decrease | K changes |
Effect of Changing Pressure (gas-phase equilibria)
Pressure changes affect equilibria only when there are different numbers of moles of gas on each side (Δngas ≠ 0).
Fig 7.2 - Effect of Pressure on 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)Diagram: Two containers side by side. Left container (low pressure): 3 moles of gas total on the left side, 2 moles on the right. Right container (high pressure): arrows show equilibrium shifting right (towards fewer gas moles, i.e. 2 moles of SO₃). Caption: "Increasing pressure shifts equilibrium towards the side with fewer moles of gas."
| Change | Shift Direction | Condition |
|---|---|---|
| Increase pressure | Towards fewer moles of gas | Only if Δn(gas) ≠ 0 |
| Decrease pressure | Towards more moles of gas | Only if Δn(gas) ≠ 0 |
| Change pressure | No shift | If equal moles of gas on both sides |
Effect of a Catalyst
Catalyst and Equilibrium A catalyst speeds up both the forward and reverse reactions equally. It does not shift the position of equilibrium and does not change the value of K. Its only role is to allow equilibrium to be reached more quickly.Summary - Effect on the Equilibrium Constant K
| Condition Changed | Position of Equilibrium Shifts? | Value of K Changes? |
|---|---|---|
| Temperature increased | Yes - towards endothermic direction | Yes |
| Temperature decreased | Yes - towards exothermic direction | Yes |
| Concentration changed | Yes | No |
| Pressure changed | Yes (if Δngas ≠ 0) | No |
| Catalyst added | No | No |
Industrial Applications
The Haber Process
The Haber process synthesises ammonia (NH₃) from nitrogen (from the air) and hydrogen (from natural gas):
Equation N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹ Fig 7.3 - The Haber Process Flow DiagramDiagram: Flow chart showing: N₂ (from fractional distillation of air) + H₂ (from steam reforming of methane) → compressor (raises pressure to ~200 atm) → reaction vessel with iron catalyst at ~450°C → cooler (NH₃ liquefies and is removed) → unreacted N₂ and H₂ recycled back to the compressor. Yield ~15% per pass.
Industrial Conditions:- Temperature: ~450°C
- Pressure: ~200 atm
- Catalyst: Iron (with K₂O and Al₂O₃ as promoters)
- Yield: ~15% per pass; unreacted gases are recycled
| Factor | Optimal Condition for Yield | Actual Industrial Choice | Reason for Compromise |
|---|---|---|---|
| Temperature | Low (forward reaction exothermic, K increases at lower T) | ~450°C | Too low a temperature gives an unacceptably slow rate; 450°C is the economic optimum |
| Pressure | High (4 mol gas → 2 mol gas, equilibrium shifts right) | ~200 atm | Higher pressures require more expensive equipment and carry greater safety risks |
| Catalyst | - | Iron | Increases rate without affecting K; iron is cheap and readily available |
The Contact Process
The Contact process manufactures sulfuric acid (H₂SO₄) in three main stages. The key equilibrium step is the oxidation of SO₂:
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