HSC Chemistry equilibrium and acid-base reactions (Modules 5 and 6): 2026 guide

What Modules 5 and 6 ask

HSC Chemistry Modules 5 (Equilibrium and Acid Reactions) and 6 (Acid/Base Reactions) are calculation-heavy modules that together make up a substantial share of the HSC exam (NESA does not publish fixed module weightings, but recent papers consistently distribute marks across all four Year 12 modules). They test sustained chemical reasoning.

The modules connect: equilibrium provides the framework, acid-base chemistry is its most-tested application. Buffers, titrations, and indicator behaviour all derive from equilibrium principles.

Module 5: Equilibrium and Acid Reactions

The equilibrium concept

A reversible reaction reaches equilibrium when forward and reverse rates equal. Macroscopic properties (concentration, colour, pressure) stop changing, but the reaction continues dynamically.

Visual evidence of equilibrium: a brown gas (NO2) and a colourless gas (N2O4) in a sealed container reach a constant brown colour at equilibrium, not because reactions stop, but because they continue at equal rates.

The equilibrium constant Kc

For the reaction:

The equilibrium constant in terms of concentration is:

Rules:

• Pure solids and liquids are excluded (their "concentration" is constant).
• Kc is temperature-dependent only. Concentration changes shift position but not Kc.
• Large Kc means the equilibrium lies far to the right (mostly products); small Kc means far to the left (mostly reactants).

Le Chatelier's principle

If an equilibrium is disturbed, it shifts to oppose the disturbance.

Concentration changes

Add a reactant: equilibrium shifts right. Remove a product: equilibrium shifts right. (Both increase forward rate temporarily until new equilibrium reached.)

Pressure changes (gas reactions only)

Increase pressure: shifts toward the side with fewer moles of gas. Decrease pressure: shifts toward the side with more moles of gas.

Temperature changes

Increase temperature: shifts in the endothermic direction (absorbing heat). Decrease temperature: shifts in the exothermic direction.

Catalyst

No shift. Catalysts increase the rate of both forward and reverse reactions equally, reaching equilibrium faster but with the same composition.

Industrial equilibrium: the Haber process

A classic exam application. The Haber process produces ammonia:

Industrial conditions are chosen to maximise yield while keeping costs reasonable:

• High pressure (200-400 atm): shifts right (4 moles gas to 2 moles gas).
• Moderate temperature (~400°C): high temperature shifts left (the forward reaction is exothermic), but very low temperature makes the reaction too slow. The compromise is moderate temperature.
• Iron catalyst: speeds up the approach to equilibrium without affecting the equilibrium position.
• Removing ammonia: shifts right (forces continued production).

Solubility equilibria (Ksp)

For a sparingly soluble salt:

If $[\text{A}^+][\text{B}^-] > K_{sp}$: precipitation occurs.
If $[\text{A}^+][\text{B}^-] = K_{sp}$: saturated solution.
If $[\text{A}^+][\text{B}^-] < K_{sp}$: unsaturated, more solid can dissolve.

Worked example: calculating equilibrium concentrations

For the reaction $H_2 + I_2 \rightleftharpoons 2HI$ at 700K, $K_c = 49$.

If 1.0 mol of H2 and 1.0 mol of I2 are placed in a 1.0 L container, what are the equilibrium concentrations?

ICE table (Initial, Change, Equilibrium):

Taking the square root:

So at equilibrium: $[H_2] = [I_2] = 0.222$ M, $[HI] = 1.556$ M.

Module 6: Acid/Base Reactions

Brønsted-Lowry theory

Brønsted-Lowry acid: a proton donor.
Brønsted-Lowry base: a proton acceptor.

Every acid has a conjugate base; every base has a conjugate acid. For example:

HCl is the acid (donates H+); H2O is the base (accepts H+); H3O+ is the conjugate acid of H2O; Cl- is the conjugate base of HCl.

pH and pOH

For dilute aqueous solutions at 25°C:

(For pure water, $[H^+] = [OH^-] = 10^{-7}$, so pH = pOH = 7.)

Strong vs weak acids and bases

Strong acids dissociate completely (HCl, HNO3, H2SO4, HClO4). For a 0.10 M HCl, $[H^+] = 0.10$ M, pH = 1.

Weak acids dissociate only partially. The dissociation constant Ka measures the extent of dissociation. For acetic acid (CH3COOH), Ka ≈ $1.8 \times 10^{-5}$. A 0.10 M acetic acid solution has $[H^+]$ much less than 0.10 M.

Worked example: pH of a weak acid

Calculate the pH of 0.10 M acetic acid (Ka = $1.8 \times 10^{-5}$).

ICE table for $CH_3COOH \rightleftharpoons CH_3COO^- + H^+$:

Buffers

A buffer resists pH change. Typically a weak acid plus its conjugate base.

Henderson-Hasselbalch equation:

When the concentrations of the weak acid and its conjugate base are equal, $pH = pK_a$.

A buffer is most effective within ±1 pH unit of its pKa.

Example: the blood buffer. Carbonic acid (H2CO3) and bicarbonate (HCO3-) keep blood pH at ~7.4. The pKa of carbonic acid is about 6.4, so the ratio $[HCO_3^-]/[H_2CO_3]$ is approximately 20:1 at blood pH.

Titrations

A titration delivers one reagent (the titrant) to another (the analyte) until the equivalence point - when stoichiometrically equivalent moles have been mixed.

Titration curves plot pH against volume of titrant added. Four shapes to recognise:

1. Strong acid / strong base: pH 7 at equivalence. Steep transition at equivalence.
2. Weak acid / strong base: pH > 7 at equivalence (conjugate base is basic). Gentle transition.
3. Weak base / strong acid: pH < 7 at equivalence (conjugate acid is acidic). Gentle transition.
4. Strong acid / weak base: pH < 7 at equivalence.

Indicators are weak acids that change colour at their pKin. Choose an indicator whose colour change range straddles the equivalence pH.

• Methyl orange: changes around pH 3-5 (use for strong-acid / weak-base titrations).
• Bromothymol blue: changes around pH 6-8 (use for strong-acid / strong-base).
• Phenolphthalein: changes around pH 8-10 (use for weak-acid / strong-base).

Common HSC Modules 5-6 traps

Ignoring temperature in Kc

Kc only changes with temperature, not with concentration or pressure.

Forgetting that pure solids/liquids are excluded from Kc

Common in heterogeneous equilibrium questions.

Confusing Le Chatelier shifts

Concentration changes shift position; temperature changes shift both position and Kc.

Approximating x in weak acid calculations when you shouldn't

The approximation $0.10 - x \approx 0.10$ is valid only if x is less than 5% of 0.10. Always check.

Choosing the wrong indicator

Markers reward indicator selection that matches the equivalence pH.

How Modules 5 and 6 are examined

In the HSC Chemistry exam:

• Multiple choice. 6-8 questions on these modules. Quick equilibrium shifts, pH estimates, indicator choices.
• Section II short questions (3-5 marks). Single-step calculations (pH, Kc, percent dissociation).
• Section II extended response (6-9 marks). Multi-step problems with ICE tables. Industrial equilibrium evaluations. Titration interpretation.

Check your knowledge

A mix of definitional, calculation/explanation, and exam-style multi-part questions covering this topic. Aim to answer all under exam conditions, then check against the solutions block.

1. Define the term dynamic equilibrium and explain why a sealed flask containing $N_2O_4$ and $NO_2$ retains a constant brown colour at equilibrium even though molecules are still reacting. (3 marks)
2. For the contact process step $2SO_{2(g)} + O_{2(g)} \rightleftharpoons 2SO_{3(g)}$, $\Delta H = -197$ kJ mol$^{-1}$. State, with reasoning, the effect on (a) the equilibrium position and (b) the value of $K_c$ of (i) increasing the pressure, (ii) increasing the temperature, (iii) adding a $V_2O_5$ catalyst. (6 marks)
3. A 2.00 L vessel contains $0.40$ mol $PCl_5$, $0.20$ mol $PCl_3$ and $0.30$ mol $Cl_2$ at equilibrium for $PCl_{5(g)} \rightleftharpoons PCl_{3(g)} + Cl_{2(g)}$. Calculate $K_c$ to three significant figures. (3 marks)
4. The solubility of barium sulfate, found as a contaminant in a sample taken from a NSW industrial drainage canal, is $1.05 \times 10^{-5}$ mol L$^{-1}$ at 25 degrees C. (a) Write the dissolution equation and the $K_{sp}$ expression. (b) Calculate $K_{sp}$. (c) Predict whether a precipitate will form when 50.0 mL of $1.0 \times 10^{-4}$ M $Ba(NO_3)_2$ is mixed with 50.0 mL of $1.0 \times 10^{-4}$ M $Na_2SO_4$. (6 marks)
5. 25.0 mL of 0.150 M propanoic acid ($K_a = 1.35 \times 10^{-5}$) is titrated against 0.150 M NaOH. (a, 2) Calculate the initial pH. (b, 2) Calculate the pH at half-equivalence. (c, 3) Calculate the pH at equivalence. (d, 2) Select an appropriate indicator from methyl orange (3.1 to 4.4), bromothymol blue (6.0 to 7.6) or phenolphthalein (8.2 to 10.0), justifying your choice. (9 marks)
6. A laboratory technician at Sydney Water dilutes 10.0 mL of 0.100 M HCl to a final volume of 250.0 mL with distilled water, then takes 25.0 mL of this dilute solution and adds it to 25.0 mL of 0.0500 M NaOH. Calculate the pH of the resulting mixture, assuming the temperature is 25 degrees C and $K_w = 1.0 \times 10^{-14}$. (4 marks)
7. Compare the buffering capacity of a 1.0 L solution containing 0.50 M ethanoic acid and 0.50 M sodium ethanoate ($pK_a = 4.74$) with a 1.0 L solution containing 0.050 M ethanoic acid and 0.050 M sodium ethanoate, after each receives 5.0 mL of 1.0 M HCl. Calculate the resulting pH in each case and account for the difference. (6 marks)
8. A NSW power station scrubs sulfur dioxide flue gas using a limestone slurry to form calcium sulfite. The dissolved scrubbing equilibrium can be modelled as CO_{2(g)} + H_2O_{(l)} \rightleftharpoons H_2CO_{3(aq)} \rightleftharpoons H^+_{(aq)} + HCO_3^-_{(aq)}. Explain, using Le Chatelier's principle and your understanding of weak-acid equilibria, (a) why bicarbonate solutions are weakly basic, (b) why the carbonic acid / bicarbonate ratio in blood plasma is approximately 1:20 at pH 7.4 given $pK_a = 6.4$, and (c) how respiratory hyperventilation alters blood pH. Support your answer with calculations where relevant. (7 marks)