HSC Exam Practice

Sample response. Le Chatelier's Principle states that if a system at equilibrium is subjected to a disturbance (change in concentration, pressure, or temperature), the system will shift in the direction that partially opposes the disturbance and re-establishes equilibrium. A pressure change has no effect on equilibrium position when the number of moles of gas is equal on both sides of the balanced equation.

Marking notes. 1 mark for defining LCP (system shifts to oppose disturbance / re-establish equilibrium); 1 mark for the condition of no pressure effect (equal moles of gas on both sides); 1 mark for specifying only gaseous species are counted (or an equivalent qualification such as “solids and liquids are excluded from the mole count”). Accept any scientifically correct phrasing.

Sample response. Halving the volume doubles the pressure. N₂(g) + 3H₂(g) ⇌ 2NH₃(g): left side has 1 + 3 = 4 moles of gas; right side has 2 moles of gas. By Le Chatelier's Principle, increasing pressure shifts equilibrium toward the side with fewer moles of gas — the right side. Equilibrium shifts right; more NH₃ is produced. K_eq is unchanged.

Marking notes. 1 mark for correct mole count (4 left, 2 right); 1 mark for direction of shift (right / toward fewer gas moles); 1 mark for linking to Le Chatelier's Principle or mechanism (shift reduces pressure). Deduct 1 if student states K_eq changes.

Sample response. At constant volume, adding argon does not change the partial pressures of SO₂, O₂, or SO₃. Each reacting gas still exerts the same pressure it did before the argon was added. Because equilibrium is determined by the partial pressures (or concentrations) of the reacting species, and these are unchanged, there is no driving force to shift the equilibrium position.

Marking notes. 1 mark for stating partial pressures of reacting gases are unchanged at constant V; 1 mark for concluding no shift occurs (or: the driving force / Q vs K_eq comparison is unchanged). Accept “concentrations are unchanged” as equivalent to partial pressures unchanged.

Sample response. (a) Equilibrium position: no change — the catalyst does not shift equilibrium. (b) K_eq: no change — K_eq depends only on temperature, which is unchanged. (c) Activation energies: the catalyst lowers the activation energy of both the forward and reverse reactions equally by providing an alternative reaction pathway; neither direction is selectively favoured.

Marking notes. 1 mark for (a) — no change to equilibrium position; 1 mark for (b) — no change to K_eq; 1 mark for (c) — both E_a values lowered equally (do not award this mark if student states only forward E_a is lowered).

Sample response. Increasing temperature changes K_eq: for an exothermic forward reaction, increasing temperature shifts equilibrium left (favouring the endothermic reverse), so K_eq decreases. Increasing pressure does not change K_eq; it shifts the equilibrium position toward fewer gas moles (the product side, in this case), increasing yield at the same K_eq value. Temperature is the only variable that changes K_eq; pressure changes only the equilibrium position.

Marking notes. 1 mark for temperature — changes K_eq (decreases for exothermic forward with rising T); 1 mark for pressure — does not change K_eq; 1 mark for correct direction of pressure shift (toward product / fewer gas moles) and the distinction that this is a position change not a K_eq change.

Sample response. Immediate observation (Stage 1): the colour darkens rapidly. Compression halves the volume, instantaneously doubling the concentration of both NO₂ (brown) and N₂O₄ (colourless). The higher [NO₂] per unit volume produces a darker brown colour. No equilibrium shift has occurred yet.

After re-equilibration (Stage 2): the colour pales from the momentarily darkened state (though remains darker than the original, pre-compression baseline). By Le Chatelier's Principle, increasing pressure shifts equilibrium toward fewer moles of gas. 2NO₂(g) (2 mol left) ⇌ N₂O₄(g) (1 mol right) — equilibrium shifts right, converting brown NO₂ to colourless N₂O₄. The concentration of NO₂ decreases from its instantaneous peak, so the colour pales. K_eq is unchanged.

Marking notes. Stage 1: 1 mark for ‘colour darkens immediately’; 1 mark for explanation (concentration/density increase, not equilibrium shift). Stage 2: 1 mark for ‘colour pales from peak’ and LCP + mole count (2 NO₂ left, 1 N₂O₄ right, shifts right); 1 mark for linking NO₂ consumption (right shift) to reduced colour intensity.

Sample response (a). As pressure increases from 1 atm to 50 atm at 500°C, the equilibrium % yield of SO₃ increases, from approximately 30% at 1 atm to approximately 87% at 50 atm. The relationship is non-linear: larger gains in yield occur at lower pressures, with diminishing returns as pressure increases further. The trend is consistent with Le Chatelier's Principle prediction for a reaction that produces fewer moles of gas. 1 mark for direction (increases with pressure); 1 mark for supporting data values from the graph.

Sample response (b). 2SO₂(g) + O₂(g) ⇌ 2SO₃(g): left side 2 + 1 = 3 mol gas; right side 2 mol gas. Le Chatelier's Principle: increasing pressure shifts equilibrium toward fewer moles of gas — the right side (products). More SO₃ is formed at new equilibrium, so % yield increases. 1 mark for mole count (3 left, 2 right); 1 mark for applying LCP (shift right); 1 mark for linking to increased SO₃ yield.

Sample response (c). The forward reaction is exothermic (ΔH = −196 kJ/mol). By Le Chatelier's Principle, increasing temperature favours the endothermic direction (reverse) — equilibrium shifts left, so less SO₃ is produced and the % yield decreases. K_eq decreases as temperature rises because the ratio of equilibrium product concentrations to reactant concentrations is smaller at higher temperatures. 1 mark for identifying forward reaction is exothermic; 1 mark for LCP — rise in T shifts equilibrium left / toward endothermic direction; 1 mark for stating K_eq decreases as temperature increases.

Sample response. The statement is incorrect in two ways. First, adding a catalyst does not shift the equilibrium position. A catalyst lowers the activation energy of both the forward and reverse reactions equally, so both rates increase by the same factor and no net shift occurs. The equilibrium concentrations of SO₂, O₂, and SO₃ are unchanged. Second, the catalyst does not improve the yield of SO₃. Yield at equilibrium is determined by K_eq, which depends only on temperature. The correct role of the V₂O₅ catalyst in the Contact process is to allow equilibrium to be reached rapidly at the industrially chosen temperature (∼450°C), without changing the yield. Without the catalyst, the reaction rate at 450°C would be too slow for commercial production.

Marking notes. 1 mark for identifying the flaw: catalyst does not shift equilibrium to the right; 1 mark for mechanism: lowers E_a equally for both directions, so no net shift; 1 mark for identifying the second flaw: catalyst does not improve yield / K_eq; 1 mark for correct role of catalyst: speeds attainment of equilibrium, not yield change.

Sample response. The Haber process N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ/mol, is the industrial synthesis of ammonia and is operated at Incitec Pivot's Gibson Island plant at approximately 200 atm and 400–500°C with an iron catalyst. Each operating condition involves a trade-off between yield and rate.

Pressure (changes equilibrium position, not K_eq): The left side has 4 moles of gas (1 N₂ + 3 H₂); the right has 2 (2 NH₃). By Le Chatelier's Principle, increasing pressure shifts equilibrium right toward fewer gas moles — increasing yield. 200 atm is chosen rather than higher pressures because the engineering and capital cost of building vessels that safely contain >200 atm exceeds the additional revenue from the yield gain. This is an economic constraint: Le Chatelier says “higher is better”, but engineering says “200 atm is the practical optimum.” Pressure does not change K_eq — it changes only the equilibrium position.

Temperature (changes K_eq): The forward reaction is exothermic. Lower temperatures shift equilibrium right (more NH₃, higher K_eq) and higher temperatures shift left (less NH₃, lower K_eq). Temperature is the only factor that changes K_eq. At 300°C the yield is very high but the reaction rate is unacceptably slow — even with catalyst. At 400–500°C, the rate is commercially viable. This is the rate–yield compromise: higher temperature sacrifices yield (lower K_eq) to gain a usable rate.

Catalyst (changes neither equilibrium position nor K_eq): The iron catalyst lowers the activation energy of both the forward and reverse reactions equally. It does not preferentially accelerate the forward reaction, does not shift the equilibrium position, and does not change K_eq. Its sole industrial function is to allow equilibrium to be attained in a residence time compatible with continuous flow operation at 400–500°C. Without the catalyst, the rate would be too slow for profitability.

Assessment of Incitec Pivot conditions: The chosen conditions (200 atm, 400–500°C, Fe catalyst) represent the engineering and economic optimum: high enough pressure to achieve meaningful yield without unsustainable infrastructure cost; high enough temperature to achieve commercial rate despite the yield compromise; and a catalyst to accelerate equilibrium attainment. Neither the catalyst nor the pressure changes K_eq; temperature is the only lever for K_eq.

Marking notes. 1 mark — gas mole count and LCP applied correctly to pressure (4 left, 2 right, shifts right); 1 mark — pressure increases yield but does not change K_eq; 1 mark — identifies rate–yield trade-off for temperature (lower T = higher yield but slower rate); 1 mark — states temperature is the only factor that changes K_eq (and states direction: exothermic, higher T lowers K_eq); 1 mark — correctly describes catalyst role: lowers E_a equally for both directions, no change to position or K_eq; 1 mark — uses Incitec Pivot / Australian industrial context with at least one specific value (200 atm or 400–500°C); 1 mark — reaches an explicit evaluative judgement about why the chosen conditions represent a commercially rational compromise.