Law of Mass Action - UNSOLVED PRACTICE SET

Q7. State the Law of Mass Action. Why is it called the 'Law of Mass Action' and not the 'Law of Concentration Action'?

Q8. For the reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), write the expression for the rate of the forward reaction and the rate of the backward reaction according to the Law of Mass Action.

Q9. Why is the active mass of a pure liquid or pure solid taken as unity in equilibrium expressions? Explain with reasoning.

Q10. The Law of Mass Action applies to elementary reactions. What is an elementary reaction? Why can't it be directly applied to complex reactions?

Q11. For the equilibrium: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), write the expression for Kc. What happens to the value of Kc if the equilibrium is disturbed by adding more O₂?

Q12. Explain the difference between the rate constant (k) and the equilibrium constant (K). Are they related?

Q13. (a) State and explain the Law of Mass Action with a suitable example.

(b) For the elementary reaction: 2NO(g) + O₂(g) → 2NO₂(g), derive the rate expression using the Law of Mass Action.

(c) A student argues that the Law of Mass Action can be applied to any reaction, whether elementary or complex. Is the student correct? Explain why or why not.

Q14. (a) Explain the concept of 'active mass' in the context of the Law of Mass Action. How is active mass different from actual mass?

(b) For the reaction: CaCO₃(s) ⇌ CaO(s) + CO₂(g), write the expression for Kc and Kp. Why are the concentrations of CaCO₃(s) and CaO(s) not included?

(c) A vessel contains a mixture of N₂, H₂, and NH₃ at equilibrium. If the volume of the vessel is suddenly halved, how does the active mass of each species change? What effect does this have on the rates of forward and backward reactions?

Q15. (a) Derive the expression for the equilibrium constant Kc for the reaction:

H₂(g) + I₂(g) ⇌ 2HI(g)

starting from the Law of Mass Action.

(b) At a certain temperature, the rate constants for the forward and backward reactions are kf = 2.5 × 10⁻³ and kb = 5.0 × 10⁻⁴ respectively. Calculate the equilibrium constant Kc.

(c) What does the magnitude of Kc tell you about the position of equilibrium?

Q16. For the reaction: 2NO(g) + Cl₂(g) ⇌ 2NOCl(g)

At a certain temperature, the rate constant for the forward reaction is kf = 4.8 × 10⁻³ L²/mol²·s and for the backward reaction is kb = 1.2 × 10⁻³ L/mol·s.

(a) Write the rate expressions for the forward and backward reactions using the Law of Mass Action.

(b) Calculate the equilibrium constant Kc for this reaction.

(c) Predict whether the equilibrium lies towards the reactants or products based on the value of Kc.

Q17. In a school laboratory, a student studies the equilibrium:

N₂O₄(g) ⇌ 2NO₂(g)

At 298 K, the rate of the forward reaction is found to be:

Ratef = kf[N₂O₄] where kf = 4.5 × 10³ s⁻¹

The rate of the backward reaction is:

Rateb = kb[NO₂]² where kb = 1.2 × 10² L/mol·s

(a) Write the equilibrium constant expression Kc for this reaction.

(b) Calculate the value of Kc at 298 K.

(c) If 0.10 mol of N₂O₄ is placed in a 2.0 L flask, set up the equilibrium expression and calculate the equilibrium concentration of NO₂.

(d) What percentage of N₂O₄ has dissociated at equilibrium?

Q18. Consider the following scenario from an Indian chemical industry:

Ammonia is produced by the Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

(a) Write the rate expressions for the forward and backward reactions using the Law of Mass Action.

(b) At 500°C, kf = 2.0 × 10⁻⁴ L²/mol²·s and kb = 3.5 × 10⁻² L/mol·s. Calculate Kc.

(c) The industrial process uses a pressure of 200 atm. Explain how increasing pressure affects the active masses of the reactants and products, and how this influences the position of equilibrium.

(d) A student suggests that increasing the temperature will increase the rate of both forward and backward reactions equally. Is this reasoning complete? What other factor must be considered?