Chemistry • Year 12 • Module 5 • Lesson 5

Le Chatelier’s Principle: Concentration & Temperature

Apply Le Chatelier’s Principle to concentration-time graphs, real industrial scenarios, the blood carbonic acid buffer, and a cause-and-effect prediction — all at Band 4–5.

Apply · Band 4–5

1. Interpret the concentration–time graph

The graph below shows the concentrations of N₂O₄(g) and NO₂(g) over time for the equilibrium:

N₂O₄(g) ⇌ 2NO₂(g) ΔH = +57 kJ mol⁻¹

At time t₁, a stress is applied to the system. Study the graph then answer the questions below. 8 marks

Figure 1.1. Concentration vs time for N₂O₄(g) ⇌ 2NO₂(g) (ΔH = +57 kJ mol⁻¹) following a stress at t₁. Illustrative data — representative of published equilibrium studies on the dinitrogen tetroxide system.

1.1 Identify the stress that was applied at t₁. Justify your identification using evidence from the graph. 2 marks

1.2 Explain the direction of the equilibrium shift after t₁ using Le Chatelier’s Principle. 2 marks

1.3 The reaction is endothermic (ΔH = +57 kJ mol⁻¹). State whether K_eq changes after this stress is applied and justify your answer. 2 marks

1.4 If instead the temperature of this system had been increased at t₁, predict the shape of the concentration-time curves for both species after t₁, and state how K_eq would change. 2 marks

Stuck? Re-read Card 2 (concentration rules) and Card 4 (temperature + K_eq).

2. Interpret Haber process equilibrium data — Incitec Pivot context

Incitec Pivot operates a Haber process plant in Gibson Island, Brisbane, producing ammonia for fertilisers. The table below shows the equilibrium yield of NH₃ (as % by volume) at various temperatures and pressures for:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹

8 marks

Temperature (°C)	200 atm	300 atm	400 atm	500 atm
300	52	63	70	74
400	25	36	44	50
450	17	26	33	38
500	11	17	22	26
600	4	7	9	11

Data adapted from Atkins’ Physical Chemistry, 10th edition, Table 6C.1 (representative values).

2.1 Describe the trend in NH₃ yield as temperature increases at constant pressure (e.g. 300 atm). Use a specific figure from the table in your answer. 2 marks

2.2 Explain why increasing temperature decreases the yield of NH₃, using Le Chatelier’s Principle and the sign of ΔH. 3 marks

2.3 At 450°C and 300 atm, predict whether K_eq is greater or less than at 300°C and 300 atm. Justify your answer. 2 marks

2.4 Despite the lower equilibrium yield, Incitec Pivot operates at approximately 400–500°C rather than 300°C. Using data from the table and one additional factor not shown in the table, explain this industrial decision. 1 mark

Stuck? Re-read Card 4 temperature rules and the lesson’s Contact Process model answer (Q5).

3. Cause-and-effect chain — hyperventilation and the blood carbonic acid buffer

During hyperventilation, a person expels CO₂ faster than the body produces it. This disturbs the following two linked equilibria that maintain blood pH:

CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)

Complete the cause-and-effect chain by filling in each empty box. 5 marks (1 per step)

Cause / event	↓	Effect (fill in the blanks)
Hyperventilation: CO₂ is expelled from the lungs faster than normal.	↓
The system responds according to Le Chatelier’s Principle.	↓
H₂CO₃ is consumed by the shift.	↓
Concentration of H⁺(aq) falls.	↓
Overall outcome (so…):	↓

Stuck? Think about what “removing CO₂” means for the left side of the equilibrium, then track the chain through both arrows.

4. Predict and justify — cobalt(II) chloride humidity indicator

Silica gel packets used in pharmaceutical packaging contain cobalt(II) chloride paper that changes colour with humidity. The equilibrium is:

CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g) (forward reaction is endothermic)

A pharmaceutical technician notices the indicator paper inside a sealed package has turned deep blue after the package was accidentally left near a heat vent at 60°C for two hours. 4 marks

4.1 Predict the colour of the indicator paper when the package is returned to room temperature (20°C) and then stored in a humid environment. Justify your prediction fully, including the direction of equilibrium shift and the effect on K_eq, if any. 4 marks

Stuck? Consider two separate disturbances: (1) lowering temperature → exothermic direction favoured; (2) increasing H₂O(g) concentration → reverse favoured. Which colour is which form?

Answers — do not peek before attempting

Q1.1 — Identifying the stress (2 marks)

The stress at t₁ is addition of N₂O₄ (a reactant). Evidence: the [N₂O₄] shows a sudden vertical increase at t₁ (a discontinuity not produced by a gradual equilibrium shift), while [NO₂] is unaffected at that instant — consistent with adding N₂O₄ directly [1]. The subsequent smooth decrease in [N₂O₄] and rise in [NO₂] confirms the forward shift as the system reaches a new equilibrium [1].

Q1.2 — Direction of shift (2 marks)

The equilibrium shifts to the right (forward direction) [1]. According to Le Chatelier’s Principle, adding N₂O₄ (a reactant) increases the concentration of a reactant, raising the forward collision frequency so that forward rate > reverse rate. The system shifts right to consume some of the added N₂O₄ and produce more NO₂, partially counteracting the disturbance [1].

Q1.3 — K_eq unchanged (2 marks)

K_eq does not change [1]. The stress applied is a concentration change (addition of N₂O₄), not a temperature change. K_eq is only altered by temperature changes, because only temperature alters the activation energies of the forward and reverse reactions by different proportions. A concentration change shifts the equilibrium position but the ratio of equilibrium concentrations (K_eq expression) reaches the same numerical value at the new equilibrium [1].

Q1.4 — Temperature increase effect (2 marks)

Since the forward reaction is endothermic (ΔH = +57 kJ mol⁻¹), increasing temperature at t₁ would shift equilibrium to the right (endothermic direction). On the graph: [N₂O₄] would decrease smoothly and [NO₂] would increase smoothly to new equilibrium values (no vertical discontinuity — unlike adding a species). K_eq would increase because for an endothermic forward reaction, raising temperature shifts the ratio of products to reactants upward [1 for graph shape + direction; 1 for K_eq increases].

Q2.1 — Trend (2 marks)

As temperature increases at 300 atm, the equilibrium yield of NH₃ decreases [1]. For example, at 300 atm, yield falls from 63% at 300°C to 7% at 600°C — a decrease of 56 percentage points [1].

Q2.2 — Explanation (3 marks)

The forward reaction (N₂ + 3H₂ ⇌ 2NH₃) is exothermic (ΔH = −92 kJ mol⁻¹) [1]. According to Le Chatelier’s Principle, increasing temperature adds thermal energy to the system, which shifts the equilibrium in the endothermic direction (the reverse reaction) to partially counteract the temperature rise [1]. This shifts the equilibrium to the left, producing less NH₃ and more N₂ and H₂, so the equilibrium yield of NH₃ decreases and K_eq decreases [1].

Q2.3 — K_eq comparison (2 marks)

K_eq at 450°C is less than K_eq at 300°C [1]. Since the forward reaction is exothermic, increasing temperature shifts the equilibrium left (toward reactants), meaning the ratio [NH₃]²/([N₂][H₂]³) is smaller at 450°C — confirming K_eq decreases as temperature increases for an exothermic forward reaction [1].

Q2.4 — Industrial decision (1 mark)

Although 300°C gives a higher equilibrium yield (63%), the reaction rate at that temperature is too slow to be commercially viable — insufficient particles have enough energy to overcome the activation energy, even with a catalyst. At 400–500°C the rate is fast enough to produce NH₃ at a commercially useful rate; the lower per-pass yield (26–44%) can be compensated by recycling unreacted N₂ and H₂ [1].

Q3 — Cause-and-effect chain (5 marks)

CO₂(aq) concentration in blood plasma decreases (CO₂ is removed from the left side of the equilibrium). [1]
Le Chatelier’s Principle: the equilibrium shifts to the left to replace some of the CO₂ removed — H₂CO₃ decomposes back toward CO₂ + H₂O. [1]
H₂CO₃ concentration falls, so the right-hand equilibrium (H₂CO₃ ⇌ H⁺ + HCO₃⁻) also shifts left to partially replace H₂CO₃, consuming H⁺ and HCO₃⁻. [1]
[H⁺] in blood plasma falls — blood becomes less acidic (more alkaline). [1]
Overall: blood pH rises (respiratory alkalosis), which can cause dizziness and tingling as the altered pH disrupts nerve and muscle function. [1]

Q4.1 — Cobalt chloride colour prediction (4 marks)

The deep blue colour indicates that CoCl₂(s) (anhydrous, blue) dominates — the forward reaction was favoured at 60°C (endothermic forward reaction shifts right with increasing temperature; K_eq increases) [1].

When returned to 20°C: temperature decreases, so Le Chatelier’s Principle shifts the equilibrium in the exothermic direction (reverse), favouring the hexahydrate. The paper begins to turn pink [1]; K_eq decreases — because the forward reaction is endothermic, a decrease in temperature shifts the equilibrium toward less product (more hexahydrate, fewer anhydrous + water vapour), reducing the ratio of products to reactants and therefore decreasing K_eq [1].

When placed in a humid environment: water vapour concentration is high, adding H₂O(g) to the system (a product). Le Chatelier’s Principle shifts the equilibrium further to the left (reverse), converting more anhydrous CoCl₂ to the pink hexahydrate CoCl₂·6H₂O [1]. K_eq is unchanged by this concentration disturbance. Final colour: pink.