R2.3
How far? The extent of chemical change
5 SL · +4 HL (SL 2.3.1-2.3.4) (HL 2.3.5-2.3.7)
Unverified
← View on the mapParent topic: How much, how fast and how far?
Guiding question
Guiding questionHow can the extent of a reversible reaction be influenced?
The one big idea
A reversible reaction reaches dynamic equilibrium when forward and reverse rates are equal. The equilibrium constant K measures how far the reaction goes, and Le Châtelier's principle predicts how it responds to change. At HL, the reaction quotient Q and ΔG⦵ = −RT lnK make the predictions quantitative.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS1.4 Counting particles by mass: The moleEquilibrium constant expressions use concentrations.
- EditorialR2.2 How fast? The rate of chemical changeDynamic equilibrium is defined as equal forward and backward rates.
What rests on this
- EditorialR1.4 Entropy and spontaneity
- EditorialR3.1 Proton transfer reactions
- EditorialR3.2 Electron transfer reactions
Core concepts that must be mastered
- Dynamic equilibrium(2.3.1)Dynamic equilibrium is reached in a closed system when the forward and backward rates are equal. Both reactions continue — the system is dynamic, not static — but the net concentrations stop changing. The escalator analogy: walking up a down-escalator at matching speed looks stationary, but both movements are real. Equilibrium requires a closed system (so products cannot escape) and is reached in both physical changes (evaporation ⇌ condensation, dissolution ⇌ precipitation) and chemical changes.
- The equilibrium law(2.3.2, 2.3.3)The equilibrium law relates K to reaction stoichiometry. For a homogeneous reaction aA + bB ⇌ cC + dD: Kc = [C]^c^[D]^d^ / ([A]^a^[B]^b^) where concentrations are equilibrium values. Pure solids and liquids are omitted (their "concentration" is constant). The magnitude of K indicates the extent of reaction: K ≫ 1 means the reaction goes nearly to completion; K > 1 favours products; K = 1 means roughly equal amounts; K < 1 favours reactants; K ≪ 1 means the reaction hardly proceeds. K for the reverse reaction is 1/K for the forward reaction (at the same temperature). K is temperature dependent — it changes only with temperature, never with concentration or pressure.
- Le Châtelier's principle(2.3.4)Le Châtelier's principle predicts the qualitative effect of changes in concentration, temperature and pressure: when a system at equilibrium is disturbed, it shifts to counteract the change. Increasing reactant concentration shifts right; increasing product shifts left. Increasing temperature shifts in the endothermic direction (K changes). Increasing pressure shifts toward the side with fewer moles of gas; if moles of gas are equal on both sides, pressure has no effect; adding an inert gas at constant volume has no effect either. You must distinguish the effect on K (only temperature changes K) from the effect on equilibrium composition (concentration and pressure shift composition without changing K). Le Châtelier's principle predicts but does not explain — the explanation is that the disturbance makes Q ≠ K and the system responds until Q = K again. It can be applied to heterogeneous equilibria such as X(g) ⇌ X(aq).
- HL: The reaction quotient(2.3.5)The reaction quotient Q uses non-equilibrium concentrations. It has the same form as Kc but with current concentrations, not equilibrium ones. If Q < K, the reaction proceeds forward (more products form); if Q > K, it proceeds in reverse; if Q = K, the system is at equilibrium. You must calculate Q and determine the direction in which the reaction will proceed.
- HL: Quantitative equilibrium calculations(2.3.6)The equilibrium law quantifies equilibrium composition. You must solve problems involving K and initial/equilibrium concentrations. Set up an ICE table (Initial, Change, Equilibrium), substitute into the Kc expression, and solve. When K is very small, you can use the approximation [reactant]initial ≈ [reactant]_eqm (because the reaction barely proceeds), which avoids solving a quadratic — quadratic equations are not expected. The quadratic formula is not required; the small-K approximation ([reactant]initial ≈ [reactant]eqm) is sufficient for the assessed calculations. Only homogeneous equilibria are assessed.
- HL: Gibbs energy and equilibrium(2.3.7)K and ΔG both measure the position of equilibrium. The link is: ΔG⦵ = −RT lnK. A large K (products favoured) corresponds to a large negative ΔG⦵; a small K (reactants favoured) corresponds to a large positive ΔG⦵; K = 1 gives ΔG⦵ = 0. This equation connects the thermodynamic framework (R1.4) to the equilibrium framework — two independent ways of describing the same property. When ΔG⦵ is positive, the equilibrium mixture is mostly reactants; when negative, mostly products.
- How this sub-topic connectsEquilibrium is the "how far" companion to kinetics' "how fast" (R2.2). The dynamic nature of equilibrium connects to collision theory — both forward and reverse reactions have rates, and equality of rates defines the equilibrium state. A catalyst speeds the approach to equilibrium without changing K or the equilibrium position (it lowers Ea for both directions equally). At HL, ΔG⦵ = −RT lnK unifies this sub-topic with the entropy–Gibbs framework (R1.4). Acid–base equilibria (R3.1) are a direct application: Ka and Kb are equilibrium constants, buffer pH comes from the equilibrium expression, and the small-K approximation is essential for weak-acid calculations. Electrochemistry (R3.2) connects via ΔG⦵ = −nFE⦵cell and ΔG⦵ = −RT lnK.
Quantitative non-negotiables
- Deduce the equilibrium constant expression Kc for a homogeneous reaction.
- Determine the relationship between K for a reaction and K for its reverse.
- Apply Le Châtelier's principle to predict the effect of concentration, temperature and pressure on equilibrium composition and on K.
- At HL: calculate the reaction quotient Q and determine the direction of change.HL
- At HL: solve problems involving K and initial/equilibrium concentrations (including the small-K approximation).HL
- At HL: quadratic formula is not required; the small-K approximation ([reactant]initial ≈ [reactant]eqm) is sufficient for the assessed calculations.HL
- At HL: use ΔG⦵ = −RT lnK to link thermodynamics and equilibrium.HL
Common failure modes
M-26 — Concentration and amount are the same thing.
Confidence: likely
Why it’s wrong: Diluting a solution changes the concentration but not the amount of solute.
Correction: This matters enormously in equilibrium and titration calculations.
M-44 — At equilibrium the reaction has stopped.
Confidence: verified
Why it’s wrong: Forward and reverse reactions continue at equal, non-zero rates — dynamic, not static.
Correction: The RSC's escalator analogy: walking up a down-escalator at matching speed looks static.
M-45 — At equilibrium the concentrations of reactants and products are equal.
Confidence: verified
Why it’s wrong: Stems from confusing rate equality with concentration equality.
Correction: The concentrations are constant, not equal; their ratio is fixed by K, which can be enormous or tiny.
M-46 — Adding more reactant increases K.
Confidence: verified
Why it’s wrong: K depends only on temperature.
Correction: Adding reactant changes Q, so the system shifts to restore Q = K, but K itself is unchanged.
M-47 — A catalyst increases the yield at equilibrium.
Confidence: verified
Why it’s wrong: It speeds the approach to equilibrium in both directions equally.
Correction: The position of equilibrium and K are unaffected.
M-48 — If I add reactant, the forward reaction just keeps going until the extra is consumed.
Confidence: verified
Why it’s wrong: Treating forward and reverse reactions as isolated.
Correction: Both rates respond; the system reaches a new equilibrium with a different composition but the same K.
M-49 — Increasing pressure always shifts the equilibrium to the right.
Confidence: verified
Why it’s wrong: It shifts toward the side with fewer moles of gas.
Correction: If the moles of gas are equal on both sides, pressure has no effect. Adding an inert gas at constant volume has no effect either.
M-50 — Le Châtelier's principle explains why the shift happens.
Confidence: likely
Why it’s wrong: It predicts, it does not explain.
Correction: The explanation is that the disturbance makes Q ≠ K, and the system responds until Q = K again.
What "HL standard" actually looks like
HLHL adds the reaction quotient Q with direction-of-shift analysis, quantitative equilibrium calculations from initial concentrations (including the approximation when K is very small — quadratics not expected), and the thermodynamic link ΔG⦵ = −RT lnK connecting equilibrium to Gibbs energy (R1.4). +4 HL hours.
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
What is the likely composition of an equilibrium mixture when ΔG⦵ is positive?
- Official IB
What is the relative effect of a catalyst on the rate of the forward and backward reactions?
- Official IB
How does the value of K for the dissociation of an acid convey information about its strength?
- Official IB
Why do catalysts have no effect on the value of K or on the equilibrium composition?
- Official IB
How does the equilibrium law help us to determine the pH of a weak acid, weak base or a buffer solution?
- Official IB
How can Gibbs energy be used to explain which of the forward or backward reaction is favoured before reaching equilibrium?
- Official IB
Why does the extent of ionization of water increase as temperature increases?
- Official IB
How would you expect the equilibrium constants of strong and weak acids to compare?
- Official IB
How can we simplify calculations when equilibrium constants Ka and Kb are very small?
- Official IB
Why must buffer solutions be composed of weak acid or base conjugate systems, not of strong acids or bases?
- Official IB
How does Le Châtelier's principle enable us to interpret the behaviour of indicators and buffer solutions?
- Official IB
Secondary cells rely on electrode reactions that are reversible. What are the common features of these reactions?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.