Summary — What This Chapter Covers
- Dynamic equilibrium: forward and reverse rates are equal; concentrations stay constant; needs a closed system.
- Le Chatelier's principle: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.
- Kc uses equilibrium concentrations; Kp uses partial pressures (partial pressure = mole fraction × total pressure).
- Only temperature changes the value of K. Concentration, pressure and catalysts do not.
- Haber process and Contact process conditions are compromises between yield, rate and cost.
- Brønsted–Lowry: an acid is a proton donor, a base is a proton acceptor.
- Strong acids and bases fully dissociate in water; weak ones partially dissociate.
- Titration curves and indicator choice depend on whether the acid and alkali are strong or weak.
1. Chemical Equilibria
Reversible reactions
Dynamic equilibrium
A reversible reaction is one that can go in both directions: products can react to re-form the reactants (shown by ⇌).
Dynamic equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant. Both reactions are still happening.
A closed system is needed: no substances can enter or leave. In an open system a gaseous product could escape and the reverse reaction could not keep pace.
Le Chatelier's principle
Predicting the shift
Le Chatelier's principle: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.
| Change | Effect on position | Effect on K |
|---|
| Increase concentration of a reactant | Moves right, to use up the added reactant | None |
| Increase pressure (gases) | Moves to the side with fewer gas molecules | None |
| Increase temperature | Moves in the endothermic direction | Changes: K increases if forward is endothermic, decreases if forward is exothermic |
| Add a catalyst | No change: speeds up forward and reverse equally, equilibrium reached faster | None |
Worked Example
For N₂O₄(g) ⇌ 2NO₂(g), ΔH = +58 kJ mol⁻¹ (N₂O₄ colourless, NO₂ brown). Predict the colour change when (a) the temperature is raised, (b) the pressure is increased.
Equilibrium constants
Kc and Kp
For aA + bB ⇌ cC + dD:
Equilibrium constants
Kc = [C]c[D]d / [A]a[B]b Kp = p(C)c p(D)d / p(A)a p(B)b
- Square brackets mean concentration at equilibrium in mol dm⁻³.
- Mole fraction of a gas = moles of that gas ÷ total moles of gas.
- Partial pressure of a gas = mole fraction × total pressure. The partial pressures add up to the total pressure.
- Kp includes only gases. Units are worked out by cancelling: for N₂ + 3H₂ ⇌ 2NH₃, Kc has units (mol dm⁻³)²/(mol dm⁻³)⁴ = mol⁻² dm⁶ and Kp has units Pa⁻².
- A large K means the equilibrium lies well to the right.