Chemical Equilibria: Le Chatelier's Principle and the Equilibrium Constant Kc

This guide covers Chemical equilibria, Le Chatelier’s principle and Kc, part of the Physical chemistry section of OxfordAQA International AS and A-level Chemistry 9620. It is AS-level content, assessed in Unit 2 – the specification is modular, assessed by units rather than papers, and this content sits within Unit 2.

Before studying this

At International GCSE you met reversible reactions and dynamic equilibrium qualitatively — see Quantitative Chemistry for the mole-calculation skills this page assumes, and your GCSE course’s treatment of reversible reactions for the qualitative picture of Le Chatelier’s principle. This page keeps that qualitative picture but adds the quantitative side: writing an equilibrium-constant expression and using it in calculations, not just describing a shift in words.

Dynamic equilibrium

A reversible reaction can proceed in both the forward and reverse directions. In a closed system, a reversible reaction reaches dynamic equilibrium: the forward and reverse reactions continue happening, but at exactly equal rates, so the concentrations of reactants and products stop changing. The reaction hasn’t stopped — the two directions have simply balanced.

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.

Applying it to each type of change:

• Concentration: increasing a reactant’s concentration shifts the position of equilibrium towards the products; increasing a product’s concentration shifts it back towards the reactants.
• Pressure (gas equilibria only): increasing pressure shifts the position of equilibrium towards the side with fewer moles of gas.
• Temperature: increasing temperature shifts the position of equilibrium in the endothermic direction, since that absorbs some of the added heat.
• Catalyst: speeds up the forward and reverse reactions equally, so it has no effect on the position of equilibrium — only on how quickly equilibrium is reached.

Industrial compromise: rate versus yield

Both the Haber process (N₂ + 3H₂ ⇌ 2NH₃) and the Contact process (2SO₂ + O₂ ⇌ 2SO₃) are exothermic with fewer moles of gas on the product side, so yield alone would favour low temperature and high pressure. Yet both are run hot: the Haber process at around 450°C and 200 atm with an iron catalyst, and the Contact process at around 450°C and only 1–2 atm with a V₂O₅ catalyst.

The reasoning is always rate versus yield, plus cost. A lower temperature would raise the equilibrium yield, but the rate would become uneconomically slow, so a moderate temperature is chosen as a compromise between the two. High pressure raises the yield further, but costs far more in plant construction and running energy — which is why the Contact process, already achieving around 96% conversion even at low pressure, does not pay for the extra cost of high pressure, while the Haber process, with a much lower equilibrium yield at atmospheric pressure, does.

The catalyst does not improve the yield at all. It speeds up the forward and reverse reactions equally, so it only shortens the time taken to reach whatever equilibrium position the temperature and pressure already determine — a genuine economic benefit, but not one that operates through Le Chatelier’s principle.

The equilibrium constant, Kc

For a general reaction aA + bB ⇌ cC + dD at equilibrium, the equilibrium constant in terms of concentration is:

Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ

with each concentration (in mol dm⁻³) raised to the power of its balancing number in the equation.

Worked example. At equilibrium in a 2 dm³ sealed container, the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g) contains 0.40 mol N₂, 1.20 mol H₂ and 0.80 mol NH₃. Calculate Kc.

First convert moles to concentrations (divide by the 2 dm³ volume): [N₂] = 0.20 mol dm⁻³, [H₂] = 0.60 mol dm⁻³, [NH₃] = 0.40 mol dm⁻³.

Kc = [NH₃]² / ([N₂][H₂]³) = (0.40)² / (0.20 × 0.60³) = 0.16 / 0.0432 = 3.7 mol⁻² dm⁶

Here the total moles of gas differ on each side (4 → 2), so the concentration units don’t cancel — Kc carries units, found from the overall power difference in the expression.

What changes K, and what doesn’t

This is the most common source of confusion in this topic, because Le Chatelier’s principle describes the position of equilibrium shifting, while a separate question is whether the equilibrium constant itself changes:

• Changing concentration, pressure, or adding a catalyst shifts the position of equilibrium (or, for a catalyst, nothing at all) but leaves the value of Kc unchanged, provided temperature stays constant.
• Only changing temperature changes the value of Kc.

Common mistakes

• Treating “no effect on Kc” and “no effect on yield” as the same idea for concentration or pressure changes. Changing concentration or pressure does shift the position of equilibrium (and therefore the yield) — it just doesn’t change the numerical value of Kc.
• Forgetting to raise each concentration to the power of its balancing number. Every term in the Kc expression is raised to the power of its coefficient in the balanced equation.
• Skipping the mol → concentration conversion. Kc is defined in terms of concentration, not raw moles — always divide by the container volume first if you’re given amounts in moles.
• Assuming Kc is always unitless. It only has no units when the total moles of gas (or dissolved species) are equal on both sides of the equation.

Quick revision checklist

• Dynamic equilibrium: equal forward/reverse rates, constant concentrations, closed system required
• Le Chatelier’s principle applied to concentration, pressure, temperature and catalysts
• Writing and calculating Kc expressions, including units
• What changes Kc (temperature only) vs what changes position only (concentration, pressure, catalyst)

Related resources

• Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle (Cambridge AS & A Level) — the equivalent AS-level topic from a different awarding body
• Quantitative Chemistry (OxfordAQA IGCSE) — the GCSE-level mole calculations this page builds on
• OxfordAQA International A-level Chemistry hub

Written against OxfordAQA International AS and A-level Chemistry 9620, specification updated November 2022, https://www.oxfordaqa.com/qualifications/international-as-a-level-chemistry/, verified 2026-08-18. Always check the current specification for your examination year.