Factors Affecting Equilibrium Le Chateliers Principle - UNSOLVED PRACTICE SET

Q7. State Le Chatelier's Principle. Why is it called the 'principle of mobile equilibrium'?

Q8. For the endothermic reaction: N₂O₄(g) ⇌ 2NO₂(g) ΔH = +ve, explain the effect of:

(a) Increasing temperature

(b) Increasing pressure

Q9. A student adds concentrated HCl to the equilibrium:

Co(H₂O)₆²⁺(pink) + 4Cl⁻ ⇌ CoCl₄²⁻(blue) + 6H₂O

The solution turns more blue. Explain this observation using Le Chatelier's Principle.

Q10. Why does a catalyst not affect the position of equilibrium, even though it increases the rate of both forward and backward reactions?

Q11. In the contact process for manufacturing sulphuric acid, the reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) is exothermic. Explain why a moderate temperature of 450°C is used instead of a lower temperature.

Q12. When a bottle of cold drink is opened, bubbles of CO₂ escape vigorously. Explain this using Le Chatelier's Principle.

Q13. (a) State Le Chatelier's Principle clearly.

(b) Apply Le Chatelier's Principle to explain the effect of the following on the equilibrium:

2NO₂(g) ⇌ N₂O₄(g) ΔH = –57.2 kJ

(i) Increase in temperature

(ii) Increase in pressure

(iii) Addition of N₂O₄ at constant volume

(iv) Addition of an inert gas at constant volume

(c) A student claims that adding an inert gas at constant pressure shifts the equilibrium towards the side with more moles of gas. Is the student correct? Explain.

Q14. (a) Explain the effect of concentration, pressure, temperature, and catalyst on a system at equilibrium using Le Chatelier's Principle.

(b) For the Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = –92.4 kJ, discuss the optimal conditions (temperature, pressure, catalyst) for maximum yield of ammonia. Explain the compromises made in industrial practice.

(c) Why is the removal of ammonia from the reaction mixture as it forms beneficial for the process?

Q15. (a) Consider the equilibrium in an aqueous solution:

CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)

Explain the effect of:

(i) Adding sodium acetate (CH₃COONa)

(ii) Adding NaOH

(iii) Diluting the solution

(b) A student adds a few drops of phenolphthalein to a solution of acetic acid. The solution remains colourless. When solid sodium acetate is added, the solution turns slightly pink. Explain this observation using Le Chatelier's Principle.

Q16. For the reaction: CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g) ΔH = –41.2 kJ

At 700 K, the equilibrium concentrations in a 2.0 L vessel are:

[CO] = 0.10 M, [H₂O] = 0.10 M, [CO₂] = 0.40 M, [H₂] = 0.40 M

(a) Calculate Kc for this reaction.

(b) If the temperature is increased to 800 K, predict whether Kc will increase, decrease, or remain the same. Explain using Le Chatelier's Principle.

(c) If 0.20 mol of CO is added to the system at 700 K (volume remains 2.0 L), predict the direction in which the equilibrium will shift. Calculate the reaction quotient Qc immediately after addition and compare it with Kc.

(d) After the system re-establishes equilibrium, will the new equilibrium concentration of CO₂ be greater than, less than, or equal to 0.40 M? Explain without calculating.

Q17. The decomposition of ammonium chloride is given by:

NH₄Cl(s) ⇌ NH₃(g) + HCl(g) ΔH = +176 kJ/mol

(a) Predict the effect of increasing temperature on the equilibrium position. Will more NH₄Cl decompose? Explain.

(b) At a certain temperature, the total pressure at equilibrium is 2.0 atm. Calculate Kp.

(c) If the volume of the container is doubled at constant temperature, what happens to the partial pressures of NH₃ and HCl? Does the equilibrium shift? Explain.

(d) In a school laboratory, a student heats NH₄Cl in a test tube and observes white fumes at the mouth of the tube. When the tube is cooled, the white fumes disappear and solid NH₄Cl reforms. Explain this observation using Le Chatelier's Principle.

Q18. In India, the production of sulphuric acid via the contact process is vital for many industries:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = –198 kJ

(a) A factory uses a 2:1 molar ratio of SO₂ to O₂. Explain why excess O₂ is not used industrially, despite Le Chatelier's Principle suggesting it would increase SO₃ yield.

(b) The reaction is carried out at 450°C and 1-2 atm pressure in the presence of V₂O₅ catalyst. Explain:

(i) Why a moderate temperature is chosen (not too low, not too high)

(ii) Why the pressure is only slightly above atmospheric (not very high)

(iii) Why a catalyst is essential

(c) In a particular reactor at 450°C, the equilibrium partial pressures are: P_SO₂ = 0.50 atm, P_O₂ = 0.25 atm, P_SO₃ = 4.50 atm. Calculate Kp.

(d) The factory manager considers increasing the pressure to 10 atm to increase SO₃ yield. Calculate the new equilibrium partial pressure of SO₃ if the initial ratio is maintained, and comment on whether this pressure increase is economically justified.