Equilibrium

Syllabus Unit: Physical Chemistry

Why it matters

Chapter overview

This topic covers dynamic chemical equilibrium — the law of mass action and the equilibrium constants Kc/Kp, how Le Chatelier's principle predicts the shift when an equilibrium system is disturbed, ionic equilibrium in acids and bases via pH, and the solubility/buffer equilibria that extend the same mass-action logic to sparingly soluble salts and pH-resistant solutions.

Chemistry Syllabus Physical Chemistry ~25% weightage

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K = 10¹⁰ (Product-favoured)

K ≈ 1 (Balanced)

K = 10⁻¹⁰ (Reactant-favoured)

1Law of Mass Action & Kc/Kp

The law of mass action states that at equilibrium, the ratio of product to reactant concentrations, each raised to its stoichiometric power, is a constant — Kc or Kp — characteristic of that reaction at a given temperature.

Learning Objectives

State the law of mass action and write the Kc expression for a given balanced equilibrium.
Distinguish Kc (concentration-based) from Kp (partial-pressure-based) and convert between them using Kp = Kc(RT)^Δn.
Explain why pure solids and liquids are omitted from the equilibrium constant expression.
Interpret the magnitude of K: K>>1 favours products, K<<1 favours reactants.

Key Terms

Law of Mass ActionThe rate of a reaction is proportional to the product of the active masses (concentrations) of the reactants, each raised to its stoichiometric coefficient.

Equilibrium Constant (Kc)The constant ratio [products]^coeff / [reactants]^coeff at equilibrium, at a fixed temperature, expressed in molar concentrations.

KpThe equilibrium constant expressed in partial pressures of gaseous species, related to Kc by Kp = Kc(RT)^Δn.

Dynamic EquilibriumThe state where forward and reverse reaction rates are equal, so concentrations stay constant over time even though both reactions keep occurring.

Chemical equilibrium is dynamic, not static: the forward and reverse reactions never stop, but their rates become exactly equal, so all concentrations stay constant.

aA + bB ⇌ cC + dD Kc = [C]^c [D]^d / [A]^a [B]^b

Kp and the Kc–Kp relationship

Kp = Kc (RT)^Δn, where Δn = (moles of gaseous products) − (moles of gaseous reactants)

If Δn = 0, then Kp = Kc.
Pure solids and liquids have a fixed 'active mass' (constant density), so they're omitted from the expression, e.g. for CaCO3(s) ⇌ CaO(s) + CO2(g), Kc = [CO2].

Common MistakeR must be in L·atm·K⁻¹·mol⁻¹ (0.0821) when using Kp = Kc(RT)^Δn with concentrations in mol/L and pressure in atm — a frequent unit-mismatch trap.

Interpreting K

K >> 1 (e.g. 10¹⁰) — reaction goes almost to completion, strongly product-favoured.
K << 1 (e.g. 10⁻¹⁰) — reaction barely proceeds, strongly reactant-favoured.
K ≈ 1 — significant amounts of both reactants and products present at equilibrium.

Exam TipK is independent of initial concentrations and of the presence of a catalyst — a catalyst speeds up how fast equilibrium is reached but never changes the value of K itself.