Q4. Using Ksp = 5.5 × 10−6 for Ca(OH)2 (NCERT Table 6.9), use the ideal concentration model to estimate its molar solubility and the pH of its saturated solution (lime water) at 298 K.
4S3 = 5.5 × 10−6, so S ≈ 1.11 × 10−2 M and [OH−] = 2S ≈ 2.2 × 10−2 M. pOH ≈ 1.65, so pH ≈ 12.35 in this ideal model. Measured saturated-lime-water pH is typically around 12.4–12.5; activity effects and the chosen constant matter. NCERT Table 6.5 (page 195) lists lime water at 10.5, which is too low for a saturated solution.
Q5. 50 mL of 0.010 M Pb(NO3)2 is mixed with 50 mL of 0.010 M NaCl. Does PbCl2 (Ksp = 1.6 × 10−5) precipitate?
After mixing, [Pb2+] = [Cl−] = 0.0050 M. Qsp = [Pb2+][Cl−]2 = 0.0050 × (0.0050)2 = 1.25 × 10−7 < 1.6 × 10−5, so no precipitate forms.
Q6. Find the molar solubility of BaSO4 (Ksp = 1.1 × 10−10) in 0.010 M Na2SO4, and explain why excess sulphate is added when barium is estimated gravimetrically.
S = 1.1 × 10−10/0.010 = 1.1 × 10−8 M, about 1000 times less than in water (1.05 × 10−5 M). The common ion SO42− makes the precipitation of Ba2+ practically complete, so very little barium is lost in the solution.
Q7. Why does calcium phosphate dissolve more readily in dilute acid than in water?
Phosphate is the anion of a weak acid. Added H+ turns PO43− into HPO42− and H2PO4−, lowering [PO43−], so Qsp falls below Ksp and more salt dissolves (NCERT Section 6.13.2, page 206: solubility rises as the pH falls).
Q8. How does passing HCl gas into a saturated solution of common salt purify it?
The extra Cl− raises the ionic product [Na+][Cl−] (strictly, the product of activities) above its saturation value, so pure NaCl crystallises out, while the more soluble impurities, such as sodium and magnesium sulphates, stay in solution.