Questions: Le Chatelier's Principle and Equilibrium Shifts

Question 1

The Haber process produces ammonia via N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92 kJ/mol). An industrial chemist wants to maximize equilibrium yield of NH₃. Which temperature strategy does Le Chatelier's principle recommend?

• A. Higher temperature — faster kinetics means more NH₃ produced per unit time
• B. Lower temperature — since the reaction is exothermic, heat is a product; removing heat (lowering temperature) shifts equilibrium toward NH₃ and increases K
• C. Higher temperature — K is always larger at higher temperatures for industrial reactions
• D. Temperature doesn't matter — only pressure and catalyst affect ammonia yield

Answer: (B)

For an exothermic reaction, heat can be treated as a product. Raising temperature adds 'heat' to the product side, shifting equilibrium leftward and decreasing K. Lowering temperature removes heat from the product side, shifting equilibrium rightward toward NH₃ and increasing K. Option A is the classic kinetics/thermodynamics confusion: higher temperature improves reaction rate but worsens equilibrium position for exothermic reactions. Industrial processes must balance these competing effects — the Haber process runs at moderate temperatures (~450°C) as a compromise, using a catalyst to speed kinetics without sacrificing too much yield.

Question 2

A sealed container holds N₂ + 3H₂ ⇌ 2NH₃ at equilibrium. Argon gas is added at constant volume. What happens to the equilibrium position?

• A. The equilibrium shifts toward NH₃ — more gas particles increase total pressure, favoring the side with fewer moles
• B. The equilibrium shifts away from NH₃ — argon dilutes the reaction mixture, shifting toward more moles of gas
• C. Nothing changes — argon does not participate in the reaction and at constant volume does not change the concentrations of any reactant or product
• D. The equilibrium shifts toward NH₃ — the inert gas increases the effective collision frequency of N₂ and H₂

Answer: (C)

At constant volume, adding an inert gas increases total pressure but does not change the concentrations of N₂, H₂, or NH₃ — each still occupies the same volume with the same number of moles as before. Since Q is expressed in terms of concentrations (or partial pressures of the reactive gases, not total pressure), Q remains equal to K and there is no shift. This is the key contrast with pressure changes that act by compressing the volume: volume compression changes all concentrations and the partial pressures of reactive gases, affecting Q. Simply adding an unreactive gas at constant volume does neither.

Question 3

Temperature is the only common equilibrium stress that changes the value of K; adding a reactant, removing a product, or increasing pressure all shift the position of equilibrium but leave K unchanged.

• T. True
• F. False

Answer: True

K is a thermodynamic quantity that depends only on temperature (via ΔG° = −RT ln K). Changes in concentration or pressure shift Q relative to K, causing the reaction to proceed forward or reverse until Q = K again — but the target K is unchanged. Only temperature alters the energy landscape of the reaction itself, changing the relative stability of reactants and products and therefore the equilibrium constant. This is why knowing whether a reaction is exothermic or endothermic is essential for temperature analysis but irrelevant for concentration and pressure analysis.

Question 4

When a stress is applied to an equilibrium system, Le Chatelier's principle predicts that the system will shift until the original conditions are fully restored.

• T. True
• F. False

Answer: False

Le Chatelier's principle states the system shifts to *partially* counteract the stress — it never fully restores the original conditions. The Explainer explicitly states: 'the word 'partially' is essential — the system never fully restores the original conditions; it reaches a new equilibrium that reduces the stress.' For example, if you add N₂ to the Haber equilibrium, the system shifts forward and consumes some of the added N₂ — but the final equilibrium has more N₂ than the original. The net effect is that the stress is dampened, not eliminated.

Question 5

Explain why Le Chatelier's principle is simply a qualitative shortcut for Q vs. K reasoning. Use a concrete example to show that both approaches give the same prediction.

Model answer: Le Chatelier's principle says 'the system shifts to counteract the stress.' Q vs. K reasoning says 'if Q < K, the reaction proceeds forward; if Q > K, it proceeds in reverse.' These are equivalent: every 'stress' changes Q relative to K, and the shift that restores Q = K is exactly the shift that counteracts the stress. Example: for N₂ + 3H₂ ⇌ 2NH₃, at equilibrium Q = K. If you add N₂, the concentration of a reactant increases. In Q terms, the denominator of Q (which contains [N₂]) grows, so Q drops below K — the reaction proceeds forward to produce more NH₃ until Q = K again. Le Chatelier says: you added a reactant, so the system shifts toward products. Same prediction, different language. The Q vs. K framework is more powerful because it is quantitative and handles cases (like adding an inert gas at constant volume) where Le Chatelier's qualitative reasoning can mislead.

Understanding that Le Chatelier's principle is a heuristic shortcut — not a separate law — prevents misapplication. When the heuristic gives ambiguous results (e.g., adding a reactant that is also a product in a complex system), falling back on Q vs. K reasoning resolves the ambiguity rigorously. Le Chatelier is useful for fast qualitative reasoning; Q vs. K is the underlying mechanism.