In simple terms
A friendly intro before the formal notes — no formulas yet.
A Balanced Tug-of-War
In a closed container, a reversible reaction never truly stops. The forward and reverse reactions keep going at exactly the same rate, so the amounts of everything stay constant even though molecules are still reacting. The equilibrium constant, Kc, is a single number that tells you where that balance sits — mostly products, mostly reactants, or a bit of both.
Picture a tug-of-war where both teams pull with exactly equal force: the rope is still under tension and both teams are still working hard, but the flag in the middle does not move. That is dynamic equilibrium — busy at the molecular level, yet unchanging on the outside. A large Kc means the flag sits deep in the products' territory; a small Kc means it barely left the reactants' side.
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Confirm the system is closed and reversible, so nothing escapes and both directions can occur.
- 2
Recognise equilibrium: forward rate = reverse rate, so concentrations become constant (not necessarily equal).
- 3
To predict a shift, apply Le Chatelier's principle — the system moves to partially oppose whatever change you impose.
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To judge the extent, write the Kc expression (products over reactants, each raised to its coefficient) and read its magnitude.
Explore the concept
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Key formulas
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Full topic notes
Formal explanation with the rigour you need for the exam.
Dynamic equilibrium and its characteristics
When a reversible reaction takes place in a closed system — one that cannot exchange matter with its surroundings — the forward reaction is fast at first and the reverse reaction is slow, because there is little product to react back. As product builds up, the reverse rate rises and the forward rate falls. Eventually the two rates become equal. From that moment the concentrations of every species stay constant, even though molecules are still reacting in both directions at the same pace. This is dynamic equilibrium: dynamic because reactions continue, equilibrium because nothing changes on the outside.
Equilibrium is reached only in a closed system (no matter enters or leaves).
At equilibrium the forward rate equals the reverse rate.
Macroscopic properties are constant — concentration, colour and (for gases) pressure no longer change.
Equilibrium is dynamic, not static: both reactions are still occurring.
Constant does not mean equal — reactant and product concentrations are usually different from each other.
The same equilibrium position is reached whether you start from reactants or from products.
Le Chatelier's principle
Le Chatelier's principle lets us predict what happens when a system at equilibrium is disturbed: the position of equilibrium shifts in the direction that partially opposes the change imposed, until a new equilibrium is established. The word 'partially' matters — the system reduces the effect of the change but never fully cancels it. Three disturbances are examinable: changing a concentration, changing pressure or volume, and changing temperature.
Only one of these disturbances changes the value of the equilibrium constant. Changing concentration or pressure shifts the position but the system re-adjusts to restore the same Kc; a catalyst leaves both position and Kc untouched. Temperature is the exception: it changes the position AND the value of Kc, because it changes the balance of the forward and reverse rates themselves.
Concentration: increasing the amount of a species shifts equilibrium AWAY from it; decreasing a species shifts equilibrium TOWARDS it. (Removing a product, for example, pulls the reaction forward.)
Pressure / volume (gases only): increasing pressure — usually by decreasing the volume — shifts equilibrium towards the side with FEWER moles of gas; decreasing pressure shifts towards MORE moles of gas. If both sides have the same number of moles of gas, changing pressure has no effect on the position.
Temperature: raising the temperature shifts equilibrium in the ENDOTHERMIC direction; lowering it shifts equilibrium in the EXOTHERMIC direction. You must know the sign of ΔH to decide which way.
Catalyst: speeds up forward and reverse reactions EQUALLY, so equilibrium is reached faster but its position — and the yield — is unchanged.
The equilibrium constant, $K_c$
Le Chatelier tells us which way a system moves; the equilibrium constant tells us how far the reaction got. For a homogeneous reaction at a fixed temperature, once equilibrium is reached the ratio of product concentrations to reactant concentrations takes a constant value called Kc. The 'c' signals that it is written in terms of molar concentrations (mol dm⁻³).
The magnitude of Kc is a direct measure of the extent of reaction. A large value means the top of the fraction dominates, so products are favoured; a small value means reactants dominate.
Square brackets mean the equilibrium molar concentration (mol dm⁻³) of a species.
Products go on top (numerator); reactants go on the bottom (denominator).
Each concentration is raised to the power of its stoichiometric coefficient from the balanced equation.
Pure solids and pure liquids are OMITTED — their concentration is constant and is absorbed into .
is constant at a given temperature; it changes value only if the temperature changes.
(roughly ): equilibrium lies far to the RIGHT; the mixture is mostly products; the reaction goes nearly to completion.
(roughly ): equilibrium lies far to the LEFT; the mixture is mostly reactants; the reaction barely proceeds.
: appreciable amounts of both reactants and products are present at equilibrium.
The reaction quotient, $Q$, versus $K_c$ (HL)
At Higher Level we often want to know the direction a reaction will move before it reaches equilibrium. The reaction quotient Q uses the same expression as Kc, but with the concentrations present at the current instant rather than at equilibrium. Comparing Q with Kc predicts what happens next: the system always moves so that Q heads towards Kc.
At HL you also calculate Kc from equilibrium concentrations, and sometimes work backwards to find a missing concentration by rearranging the expression. The method is identical to the worked example above: write the expression, substitute the equilibrium values, and evaluate — remembering to raise each concentration to its coefficient.
: too little product relative to equilibrium, so the reaction proceeds FORWARD (to the right) to make more product.
: too much product relative to equilibrium, so the reaction proceeds in REVERSE (to the left).
: the system is already at equilibrium and there is no net change.
Q changes continuously as the reaction proceeds; it stops changing once it equals Kc.
Common mistakes examiners penalise
Saying 'at equilibrium the concentrations are equal' — they are CONSTANT, not equal. This confusion is heavily penalised.
Claiming a catalyst increases yield or shifts the equilibrium — it only speeds up reaching equilibrium; position, yield and Kc are unchanged.
Getting the temperature direction wrong — increasing temperature favours the ENDOTHERMIC direction; you must use the sign of ΔH to decide, not guess.
Thinking pressure or concentration changes the value of Kc — they change the POSITION of equilibrium; only a change in temperature changes Kc.
Shifting the wrong way for pressure — higher pressure favours FEWER moles of gas; and if both sides have equal moles of gas, pressure has no effect at all.
Including solids or liquids in the Kc expression — omit pure (s) and (l); only gases and aqueous species appear.
Forgetting the stoichiometric powers — each concentration is raised to its coefficient (e.g. , ); missing a power gives the wrong Kc.
(HL) Reversing the Q vs Kc logic — means go FORWARD; means go reverse.
Model answer — marked the way our engine marks it
This is the showcase for an explanation-style question. In Paper 2 the marks are analytic: each distinct valid point scores 1 mark, typically a reasoning step (M) followed by the conclusion it supports (A). The engine credits each correct point independently and accepts equivalent wording, and error-carried-forward means a sound consequence drawn from an earlier (even mistaken) statement can still score. Study how each of the four marks below is tied to one specific idea.
Where this leads
Equilibrium ideas run through the rest of the course. Acid–base chemistry is equilibrium with Ka, Kb and Kw; solubility uses the same reasoning; industrial processes such as the Haber and Contact processes are Le Chatelier trade-offs between yield and rate. Fix the two questions of this topic — which way will it move (Le Chatelier) and how far did it get (Kc) — and every later equilibrium becomes a variation on a method you already own.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Ethanoic acid and ethanol form an ester in a reversible reaction: CH₃COOH(l) + C₂H₅OH(l) ⇌ CH₃COOC₂H₅(l) + H₂O(l). The forward reaction is used to make the ester. State and explain, using Le Chatelier's principle, TWO changes that would increase the equilibrium yield of the ester. [4]
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Change 1 — remove a product (or add a reactant). Continuously removing the ester (or the water) as it forms, OR adding an excess of one reactant such as ethanol, disturbs the balance. [M1: correct change stated] By Le Chatelier's principle the equilibrium shifts to oppose the change: removing a product shifts equilibrium to the RIGHT to replace it (adding a reactant shifts it right to use the extra reactant up), increasing the yield of ester. [A1: correct direction linked to increased yield]
Hydrogen and iodine reach equilibrium: H₂(g) + I₂(g) ⇌ 2HI(g). (a) Write the expression for Kc. (b) At a certain temperature the equilibrium concentrations are [H₂] = 0.20 mol dm⁻³, [I₂] = 0.20 mol dm⁻³ and [HI] = 1.60 mol dm⁻³. Calculate Kc and comment on the extent of reaction. [4]
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(a) Kc expression. — product on top, squared because its coefficient is 2; both reactants on the bottom. [M1: correct expression]
For the Haber process, N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH < 0. Explain the effect of (i) increasing the pressure and (ii) increasing the temperature on the position of equilibrium and on the yield of ammonia. [4]
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Model answer.
How it all connects
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Tap a linked idea to see how it connects back to the main topic — that connection is what examiners reward.
Glossary
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Quick check
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Revision flashcards
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Dynamic equilibrium
In a closed system, the state reached when the forward and reverse reactions occur at equal rates, so the concentrations of reactants and products remain constant. The reaction has not stopped — both directions continue.
Key takeaways
Review these before you close the topic — retrieval beats re-reading.
- ✓
Equilibrium is reached only in a closed system (no matter enters or leaves).
- ✓
At equilibrium the forward rate equals the reverse rate.
- ✓
Macroscopic properties are constant — concentration, colour and (for gases) pressure no longer change.
- ✓
Equilibrium is dynamic, not static: both reactions are still occurring.
- ✓
Constant does not mean equal — reactant and product concentrations are usually different from each other.
- ✓
The same equilibrium position is reached whether you start from reactants or from products.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Get a Paper 2 answer marked: apply Le Chatelier's principle and reason with Kc, with full working
Get a Paper 2 answer marked: apply Le Chatelier's principle and reason with Kc, with full working
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Checkpoint
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Before you move on: do Get a Paper 2 answer marked: apply Le Chatelier's principle and reason with Kc, with full working on paper, snap a photo, and get examiner-style feedback on exactly where you win and lose marks.