Solubility Product and Common Ion Effect - UNSOLVED PRACTICE SET

Q7. Define solubility product (Ksp). Write the Ksp expression for Ca₃(PO₄)₂.

Q8. Explain the common ion effect with a suitable example. How does it affect the solubility of a sparingly soluble salt?

Q9. The Ksp of AgCl at 25°C is 1.8 × 10⁻¹⁰. Calculate its solubility in pure water in mol/L.

Q10. Distinguish between solubility and solubility product. Are they the same? Explain.

Q11. Predict whether a precipitate will form when equal volumes of 0.01 M AgNO₃ and 0.01 M NaCl are mixed. (Ksp of AgCl = 1.8 × 10⁻¹⁰)

Q12. Explain why the solubility of Ca(OH)₂ decreases when NaOH is added to its saturated solution.

Q13. (a) Define solubility product and derive the relationship between solubility product (Ksp) and solubility (s) for:

(i) AB type salt (e.g., AgCl)

(ii) A₂B type salt (e.g., Ag₂CrO₄)

(iii) AB₂ type salt (e.g., CaF₂)

(b) The Ksp of CaF₂ is 3.9 × 10⁻¹¹. Calculate its solubility in pure water.

(c) Explain how the common ion effect is used in qualitative analysis to separate metal ions.

Q14. (a) State and explain the common ion effect with reference to the solubility of AgCl in:

(i) Pure water

(ii) 0.1 M NaCl solution

(iii) 0.1 M AgNO₃ solution

(b) Calculate the solubility of AgCl in 0.1 M NaCl solution. (Ksp of AgCl = 1.8 × 10⁻¹⁰)

(c) Compare this solubility with that in pure water and explain the difference.

Q15. (a) Explain the conditions under which precipitation occurs. Distinguish between saturated, unsaturated, and supersaturated solutions using the concepts of ionic product (Q) and Ksp.

(b) A solution contains 0.01 M Mg²⁺ ions and 0.01 M OH⁻ ions. Will Mg(OH)₂ precipitate? (Ksp of Mg(OH)₂ = 1.8 × 10⁻¹¹)

(c) What is the minimum concentration of OH⁻ ions needed to start precipitation of Mg(OH)₂ from a 0.01 M Mg²⁺ solution?

Q16. Calculate the solubility product (Ksp) for the following salts given their solubilities:

(a) AgCl: solubility = 1.34 × 10⁻⁵ mol/L

(b) CaF₂: solubility = 2.15 × 10⁻⁴ mol/L

(c) Ag₂CrO₄: solubility = 1.3 × 10⁻⁴ mol/L

(d) Al(OH)₃: solubility = 2.9 × 10⁻⁹ mol/L

Also write the dissociation equation for each salt.

Q17. The Ksp values at 25°C are:

AgCl = 1.8 × 10⁻¹⁰, AgBr = 5.0 × 10⁻¹³, AgI = 8.3 × 10⁻¹⁷

(a) Calculate the solubility of each salt in pure water.

(b) Arrange the salts in order of increasing solubility. Does this order match the order of Ksp values? Explain why or why not.

(c) A solution contains 0.01 M Cl⁻, 0.01 M Br⁻, and 0.01 M I⁻. If AgNO₃ is added slowly, which salt will precipitate first? Calculate the [Ag⁺] required to start precipitation of each salt.

(d) What is the [Br⁻] remaining in solution when AgCl just begins to precipitate?

Q18. In India, hard water containing Ca²⁺ and Mg²⁺ ions is a common problem:

(a) Explain how washing soda (Na₂CO₃) is used to remove hardness. Write the relevant solubility product expressions and explain the common ion effect involved.

(b) The Ksp of CaCO₃ is 4.8 × 10⁻⁹. Calculate the solubility of CaCO₃ in:

(i) Pure water

(ii) 0.05 M Na₂CO₃ solution

(c) A water sample contains 0.001 M Ca²⁺ ions. What minimum concentration of CO₃²⁻ is needed to start precipitation of CaCO₃?

(d) Kidney stones often contain calcium oxalate (CaC₂O₄, Ksp = 2.3 × 10⁻⁹). A patient has 0.005 M Ca²⁺ in their urine. Calculate the maximum concentration of oxalate ions (C₂O₄²⁻) that can be present without causing precipitation. Suggest dietary advice based on this calculation.