IB DP Chemistry Reactivity 2.3: How far? The extent of chemical change – Revision Notes

These notes condense Reactivity 2.3, How far? The extent of chemical change from the IB Diploma Programme Chemistry guide (February 2023, first assessment 2025) for final revision. They cover understandings Reactivity 2.3.1 to 2.3.7: 2.3.1 to 2.3.4 are for SL and HL, and 2.3.5 to 2.3.7 are HL only. The study guide explains each idea in full with worked examples.

Course links: IB DP chemistry hub, printable checklist, and the practice questions for this unit. Our free diagnostics help you find weak topics.

Key definitions

• Reversible reaction: one in which products can react to re-form reactants (shown by ⇌).
• Closed system: no matter is exchanged with the surroundings. Equilibrium needs one.
• Dynamic equilibrium: the state in which forward and backward rates are equal, so concentrations stay constant while both reactions carry on.
• Homogeneous equilibrium: every species is in the same phase.
• Heterogeneous equilibrium: species are in more than one phase, for example X(g) ⇌ X(aq).
• Equilibrium constant, K: the value of the equilibrium expression at equilibrium, at a stated temperature.
• Le Châtelier’s principle: a disturbed equilibrium shifts its position to partly oppose the disturbance.
• Reaction quotient, Q (HL only): the equilibrium expression evaluated with concentrations at any moment.

Reactivity 2.3.1 Features of equilibrium

Feature	At equilibrium
Forward and backward rates	Equal, and not zero
Concentrations	Constant (not necessarily equal)
Observable properties	Colour, pressure and density constant
System	Closed
Approach	Can be reached from either side

A physical example: a sealed flask part-filled with a volatile liquid at constant temperature. Evaporation and condensation rates become equal and the vapour pressure stays constant. A chemical example: a sealed mixture of a reversible gas reaction once its composition stops changing.

Reactivity 2.3.2 Writing K

For aA + bB ⇌ cC + dD (homogeneous):

K = [C]ᶜ[D]ᵈ / ([A]ᵃ[B]ᵇ)

Method in steps:

1. Check the equation is balanced and every species is in the same phase.
2. Products on top, reactants underneath.
3. Each concentration raised to the power of its coefficient.
4. Multiply terms; never add them.

Small reminder: for 2N₂O(g) + 3O₂(g) ⇌ 4NO₂(g), K = [NO₂]⁴ / ([N₂O]²[O₂]³).

Reactivity 2.3.3 Size of K and reversed equations

K	Meaning
K ≪ 1	Hardly any reaction; reactants dominate
K < 1	More reactants than products
K = 1	Neither side favoured
K > 1	More products than reactants
K ≫ 1	Reaction nearly complete; products dominate

• K(reverse) = 1 / K(forward) at the same temperature.
• K changes only with temperature.
• A large K says nothing about rate.

Small reminder: K = 0.25 forward means K = 4.0 for the reverse.

Reactivity 2.3.4 Le Châtelier’s principle

Disturbance	Position moves	K
Increase [reactant]	Right	Same
Increase [product]	Left	Same
Remove a product	Right	Same
Increase pressure	To fewer gas molecules	Same
Decrease pressure	To more gas molecules	Same
Raise temperature	Endothermic direction	Changes
Lower temperature	Exothermic direction	Changes
Add catalyst	No shift; equilibrium reached faster	Same

Temperature and K:

Temperature change	Forward reaction exothermic	Forward reaction endothermic
Heating	Smaller K, less product	Larger K, more product
Cooling	Larger K, more product	Smaller K, less product

If both sides have equal numbers of gas molecules, pressure has no effect on position.

Heterogeneous example: NH₃(g) ⇌ NH₃(aq). Raising the partial pressure of ammonia above the solution pushes more into solution. Dissolving ammonia is exothermic, so warming the solution drives gas out.

Method in steps for a “predict and explain” answer:

1. Name the change and the direction that opposes it.
2. State the shift (left or right) and what happens to a named species.
3. Give the reason (fewer gas molecules, endothermic direction, replacing a removed species).
4. Say whether K changes.

Reactivity 2.3.5 Q and direction (HL only)

Q compared with K	Net change
Q < K	Forward, making more products
Q = K	None: at equilibrium
Q > K	Backward, making more reactants

Small reminder: Q is built from the same expression as K. If K = 5.0 and a mixture gives Q = 0.80, the mixture has too few products, so the forward reaction runs until Q rises to 5.0. Q can only be compared with K measured at the same temperature.

Reactivity 2.3.6 Equilibrium calculations (HL only)

Method in steps (ICE table):

1. Write the balanced equation and the K expression.
2. Initial row: starting amounts or concentrations.
3. Change row: use the mole ratio (e.g. −x, −2x, +x).
4. Equilibrium row; convert moles to mol dm⁻³ by dividing by volume.
5. Substitute into K and solve.

Small reminder: SO₂Cl₂(g) ⇌ SO₂(g) + Cl₂(g). 0.50 mol SO₂Cl₂ is heated in a 2.0 dm³ flask; at equilibrium 0.10 mol Cl₂ is present. So 0.10 mol SO₂ also forms and 0.40 mol SO₂Cl₂ remains. Concentrations: [SO₂Cl₂] = 0.20, [SO₂] = [Cl₂] = 0.050 mol dm⁻³. K = (0.050 × 0.050) / 0.20 = 0.0125.

Shortcuts that avoid quadratics:

• Perfect square: if the expression becomes x² / (c − x)², take square roots of both sides.
• Small K: if K is very small, take [reactant]eqm ≈ [reactant]initial, then check the change is tiny compared with the starting value.

Only homogeneous equilibria are assessed here, and the guide states quadratic equations are not expected.

Reactivity 2.3.7 K and Gibbs energy (HL only)

ΔG⦵ = −RT lnK R = 8.31 J K⁻¹ mol⁻¹, T in K, ΔG⦵ in J mol⁻¹

ΔG⦵	K	Position
< 0	> 1	Products favoured
= 0	= 1	Neither favoured
> 0	< 1	Reactants favoured

Rearranged: K = e^(−ΔG⦵/RT).

Small reminder: ΔG⦵ = +10.0 kJ mol⁻¹ at 400 K gives lnK = −10 000 / (8.31 × 400) = −3.01, so K = 0.0494. The positive ΔG⦵ and K below 1 both say reactants are favoured.

Must-know distinctions

• Equal rates vs equal concentrations: only the rates are equal.
• Where the equilibrium lies vs the number K: a change of concentration or pressure moves the mixture, but K itself responds to temperature alone.
• How far (K, ΔG⦵) vs how fast (Ea, rate): a catalyst affects only how fast.
• Q (any moment) vs K (equilibrium only).
• Moles vs concentration: K uses mol dm⁻³.

Quick self-test

1. State two features of a chemical system at dynamic equilibrium.
2. Deduce K for CS₂(g) + 3Cl₂(g) ⇌ CCl₄(g) + S₂Cl₂(g).
3. K for a reaction is 6.4 × 10³ at 450 K. State K for the reverse reaction at 450 K.
4. A reaction has K = 1.0 × 10⁻¹² at 298 K. Describe the equilibrium mixture.
5. For C₂H₆(g) ⇌ C₂H₄(g) + H₂(g), ΔH positive, state the effect of increasing the pressure on the yield of ethene and on K.
6. For the same reaction, state the effect of raising the temperature on K.
7. State the effect of a catalyst on the composition of an equilibrium mixture.
8. (HL only) For the reaction in question 5, K = 0.050 at a set temperature. A mixture has [C₂H₆] = 0.80, [C₂H₄] = 0.10 and [H₂] = 0.20 mol dm⁻³. Calculate Q and state the direction of net change.
9. (HL only) Br₂(g) ⇌ 2Br(g), K = 4.0 × 10⁻⁶ at a set temperature. [Br₂] starts at 0.100 mol dm⁻³. Calculate [Br] at equilibrium.
10. (HL only) K = 50 at 600 K. Calculate ΔG⦵.
11. (HL only) State ΔG⦵ for a reaction with K = 1.

Answers

1. Any two: forward and backward rates are equal; concentrations stay constant; both reactions continue; it is in a closed system.
2. K = [CCl₄][S₂Cl₂] / ([CS₂][Cl₂]³)
3. 1 / (6.4 × 10³) = 1.6 × 10⁻⁴
4. Almost entirely reactants; hardly any products form.
5. Yield of ethene falls (position moves to the side with fewer gas molecules, the left); K unchanged.
6. Forward is endothermic, so K increases.
7. No effect; equilibrium is reached sooner.
8. Q = (0.10 × 0.20) / 0.80 = 0.025. Q < K, so net forward reaction.
9. [Br]² ≈ 4.0 × 10⁻⁶ × 0.100 = 4.0 × 10⁻⁷, so [Br] = 6.3 × 10⁻⁴ mol dm⁻³.
10. ΔG⦵ = −8.31 × 600 × ln 50 = −19 505 J mol⁻¹ = −19.5 kJ mol⁻¹.
11. Zero, because ln 1 = 0.

Where marks are usually lost

• Writing “the concentrations of reactants and products are equal” as a feature of equilibrium.
• Adding terms in a K expression, or using coefficients as multipliers rather than powers.
• Leaving out the “closed system” condition when describing equilibrium.
• Saying K changes when a reactant is added or the pressure is raised.
• Giving a shift but no reason, or no named species, in a Le Châtelier answer.
• Counting all molecules, including liquids or solutions, when judging a pressure change; count gas molecules only.
• Saying a catalyst increases the yield.
• (HL only) Substituting moles directly into K for a vessel that is not 1 dm³.
• (HL only) Reading Q > K as “moves forward”.
• (HL only) Using kJ or °C in ΔG⦵ = −RT lnK.

Official syllabus

Aligned with the IB’s DP Chemistry guide (International Baccalaureate; issued February 2023; first assessment 2025). These notes follow Reactivity 2.3, How far? The extent of chemical change, understandings 2.3.1 to 2.3.4 (SL and HL) and 2.3.5 to 2.3.7 (HL only).