Level 1 — Basic Recall (1–10)

1. What is chemical equilibrium?

2. Define “dynamic equilibrium.”

3. State Le Chatelier’s principle.

4. What is the equilibrium constant, (K_c)?

5. Write the expression for (K_c) for the reaction: (\text{aA + bB ⇌ cC + dD}).

6. What does a large (K_c) value indicate?

7. What does a small (K_c) value indicate?

8. Define homogeneous equilibrium.

9. Define heterogeneous equilibrium.

10. What is meant by “reaction quotient,” (Q)?

Level 2 — Understanding (11–20)

11. Explain why equilibrium is described as dynamic rather than static.

12. Why do catalysts not change the position of equilibrium?

13. Why are solids not included in equilibrium expressions?

14. Describe the effect of increasing concentration on equilibrium position.

15. Describe the effect of decreasing pressure on a gaseous equilibrium.

16. Explain how temperature affects the equilibrium constant.

17. Why does (K_c) remain unchanged when concentrations change?

18. Distinguish between endothermic and exothermic reactions in equilibrium shifts.

19. What happens when Q > K?

20. What happens when Q < K?

Level 3 — Application (21–30)

21. Write the (K_c) expression for: (N_2 + 3H_2 ⇌ 2NH_3).

22. Predict the equilibrium shift if pressure is increased in the Haber process.

23. Predict how adding a catalyst affects yield.

24. Calculate the value of Q for a given concentration set.

25. For an exothermic reaction, predict the effect of increasing temperature.

26. Given chemical concentrations at equilibrium, calculate (K_c).

27. Explain why removing product increases yield.

28. For the equilibrium: (2SO_2 + O_2 ⇌ 2SO_3), predict the effect of increasing pressure.

29. Determine whether the reaction will shift left or right when concentration changes.

30. Identify which species are omitted from the (K_c) expression in a heterogeneous system.

Level 4 — Analysis (31–40)

31. Analyse why temperature alters the equilibrium constant but pressure does not.

32. Compare large vs small (K_c) values in terms of extent of reaction.

33. Use Le Chatelier’s principle to predict shifts in multi-step reactions.

34. Discuss the industrial significance of equilibrium in ammonia synthesis.

35. Explain why equilibrium is never reached in an open system.

36. Compare equilibrium behaviour of exothermic vs endothermic reactions when heated.

37. Analyse the response of equilibrium when inert gases are added at constant pressure.

38. Determine equilibrium behaviour when volume is doubled.

39. Compare the behaviour of strong electrolytes vs weak electrolytes in equilibrium.

40. Evaluate how equilibrium principles apply to buffer solutions (qualitative).

Level 5 — Exam/Challenge (41–50)

41. Given equilibrium data for multiple trials, calculate (K_c) and comment on consistency.

42. Use an ICE table to calculate equilibrium concentrations.

43. Analyse a system with multiple equilibria (coupled equilibrium).

44. Determine the effect of temperature increase on equilibrium for given ΔH.

45. Predict changes in equilibrium yield in industrial processes using numerical data.

46. Solve a problem involving gaseous equilibrium shifting due to pressure changes.

47. Derive the relationship between (K_p) and (K_c).

48. Evaluate the thermodynamic reasoning behind equilibrium constants (ΔG° relation).

49. Use equilibrium data to calculate ΔG° for a reaction.

Provide a detailed explanation for why equilibrium constants express reaction feasibility but not reaction rate.