Level 1 – Basic Recall (1–10)

1. Define pH.

2. Write the expression for Kw (ionic product of water).

3. What is the value of Kw at 25°C?

4. What is the pH of a neutral solution at 25°C?

5. Define a strong acid with an example.

6. Define a weak acid with an example.

7. What is a buffer solution?

8. What is meant by solubility product (Ksp)?

9. Define a common ion.

10. What is meant by a sparingly soluble salt?

Level 2 – Understanding (11–20)

11. Explain the difference between strong and concentrated acid.

12. Write the expression for Ka of a weak acid HA ⇌ H⁺ + A⁻.

13. Explain why strong acids have high Ka values.

14. Why is pH scale logarithmic rather than linear?

15. Explain how a buffer solution resists pH change.

16. What is the difference between an acidic buffer and a basic buffer?

17. Explain how adding a common ion affects the solubility of a salt.

18. Why are salts like AgCl considered sparingly soluble?

19. Explain the term saturation in relation to solubility.

20. What happens when Q (ionic product) exceeds Ksp?

Level 3 – Application (21–30)

21. Calculate the pH of a 0.01 mol dm⁻³ HCl solution (assume complete dissociation).

22. Calculate the pH of a 0.001 mol dm⁻³ NaOH solution.

23. Given [H⁺] = 1 × 10⁻⁵ mol dm⁻³, find the pH.

24. Write the Ka expression and pH relation for a weak monoprotic acid.

25. Predict the effect on pH when a small amount of NaOH is added to a CH₃COOH/CH₃COO⁻ buffer.

26. Calculate [H⁺] of pure water at 25°C using Kw.

27. For the buffer made from 0.1 M CH₃COOH and 0.1 M CH₃COONa, state whether pH is < 7, = 7, or > 7.

28. Write the Ksp expression for CaF₂(s) ⇌ Ca²⁺ + 2F⁻.

29. Predict what happens when solid NaCl is added to a saturated solution of AgCl.

30. From a given Ksp value, calculate the molar solubility of a 1:1 salt (e.g. AgCl).

Level 4 – Analysis (31–40)

31. Compare the pH of 0.1 M HCl and 0.1 M CH₃COOH and explain the difference.

32. Analyse how dilution affects the degree of ionization of a weak acid.

33. Explain why blood acts as a buffer system.

34. Compare the action of a buffer made from weak acid + salt vs weak base + salt.

35. Explain how a common ion suppresses the ionization of a weak electrolyte (common ion effect).

36. Analyse the effect on solubility of CaSO₄ when some CaCl₂ is added to the solution.

37. Determine whether a precipitate will form when two ionic solutions are mixed, given Q and Ksp.

38. Explain why pH at the half-equivalence point of a weak acid–strong base titration equals pKa.

39. Compare the buffer capacities of two solutions with different acid/salt ratios.

40. Analyse how changing temperature affects Kw and therefore pH of neutral water.

Level 5 – Exam/Challenge (41–50)

41. A buffer contains 0.2 mol dm⁻³ CH₃COOH and 0.1 mol dm⁻³ CH₃COONa (Ka given).

    (a) Calculate its pH.

    (b) Predict the change in pH when a small amount of HCl is added.

42. Design an acidic buffer with pH ≈ 5 using acetic acid; state required ratio of [acid]/[salt].

43. Compare pH changes when adding the same amount of HCl to:

    (a) pure water

    (b) an acetate buffer.

44. For a saturated solution of Ag₂CrO₄, derive the expression relating Ksp to its solubility (s).

45. Given Ksp(AgCl) and concentration of Cl⁻, calculate whether Ag⁺ will begin to precipitate.

46. Explain, with calculation, how selective precipitation can separate two metal ions based on different Ksp values.

47. For a weak acid HA with known Ka, derive the approximate pH formula for a solution of concentration C (assume α ≪ 1).

48. Evaluate how adding NH₃ affects the solubility of AgCl, considering complex ion formation (qualitative if you haven’t done full complex equilibria).

49. Discuss how buffer solutions are crucial in maintaining pH in industrial processes or biological systems (pick one example and analyse in depth).

50. Given a multi-step equilibrium involving acid dissociation and solubility (e.g. weak acid + sparingly soluble salt), predict qualitatively how pH shifts will affect precipitation/dissolution.