01 · Equilibrium constant
Expressing equilibrium: the K expression
For a general reversible reaction: aA + bB ⇌ cC + dD, the equilibrium constant Kc is:
Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
Each concentration is the equilibrium value. Pure solids and pure liquids (including water in dilute aqueous solution) are excluded from the expression (their "concentration" is constant).
The magnitude of K tells you where equilibrium lies:
K ≪ 1
Reactants favoured. Little product at equilibrium.
K increasing →
K ≫ 1
Products favoured. Reaction goes nearly to completion.
K is temperature-dependent but unaffected by changes in concentration, pressure (for Kc), or catalyst. The reaction quotient Q has the same form as K but uses current (non-equilibrium) concentrations: if Q < K, the reaction proceeds forward; if Q > K, backward.
02 · Le Chatelier's principle
Predicting how equilibrium shifts
Le Chatelier's principle (1884): when a system at equilibrium is disturbed, it shifts to partially counteract the disturbance. This qualitative rule is enormously useful for industrial process design.
The thermodynamic basis: changing temperature changes K itself (via van't Hoff equation). Changing concentration or pressure does not change K — it changes Q, driving the system back toward equilibrium. See the Thermodynamics guide for the ΔG° = −RT ln K connection.
03 · Applications
Equilibrium in industrial chemistry and biology
| Process | Equation | Conditions | Equilibrium note |
|---|---|---|---|
| Haber process | N₂(g) + 3H₂(g) ⇌ 2NH₃(g) | 450 °C, 200 atm, Fe catalyst | ΔH = −92 kJ/mol. Higher pressure and lower T favour products, but low T is too slow. Compromise conditions. |
| Contact process (H₂SO₄) | 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) | 450 °C, 1–2 atm, V₂O₅ catalyst | ΔH = −197 kJ/mol. Higher pressure favours products (3 mol gas → 2 mol), but high-pressure equipment is expensive. |
| Carbonic acid in blood | CO₂(aq) + H₂O ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺ | pH 7.35–7.45, 37 °C | This equilibrium buffers blood pH. Hyperventilation (exhaling CO₂) shifts equilibrium left, raising pH (alkalosis). |
| Ester hydrolysis | RCOOR' + H₂O ⇌ RCOOH + R'OH | Acid or base catalyst | K ~ 1–4 for most esters. Excess water drives equilibrium toward hydrolysis. Base hydrolysis (saponification) is irreversible. |
The blood buffer example shows that equilibrium is not just a laboratory concept — it maintains the pH of blood within ±0.05 units around 7.40, which is essential for enzyme function. See Homeostasis for how the kidneys and lungs cooperate to maintain this equilibrium, and the Acids and bases guide for buffer chemistry.
In electrochemistry, the Nernst equation is equilibrium applied to half-cells: when Q = K, E = 0 and the cell is fully discharged. See Electrochemistry.
Equilibrium — common questions
- Equilibrium is dynamic — both forward and reverse reactions continue, but at equal rates. The equilibrium constant K expresses the ratio of product to reactant concentrations at equilibrium (each raised to its stoichiometric power). A large K (≫ 1) means products are favoured; a small K (≪ 1) means reactants are favoured. See the Reaction rates guide.
- Increasing reactant concentration → equilibrium shifts right (toward products). Increasing pressure on a gaseous equilibrium → shifts toward fewer moles of gas. Increasing temperature → shifts toward endothermic direction (removing the added heat). Decreasing temperature → shifts toward exothermic direction. Catalyst → no effect on equilibrium position, only on how quickly it is reached. Connects to thermodynamics.
- Kc uses concentrations [mol/L]; Kp uses partial pressures [Pa or atm]. They are related by Kp = Kc(RT)^Δn where Δn is the change in moles of gas. Ka is simply Kc for HA ⇌ H⁺ + A⁻. A larger Ka means a stronger acid (more dissociation). See Acids and bases.
- ΔG° = −RT ln K links equilibrium to thermodynamics. When Q < K, ΔG < 0 and the reaction proceeds forward. When Q > K, ΔG > 0 and the reaction proceeds backward. At Q = K, ΔG = 0 — equilibrium. See the Thermodynamics guide and Electrochemistry (where ΔG° = −nFE°).