The solubility product of \(BaSO_4\) in water is \(1.5 \times 10^{-9} \). The molar solubility of \(BaSO_4\) in 0.1 M solution of Ba(NO₃)₂ in- 
1. \(2.0 \times 10^{-8} M\)
2. \(0.5 \times 10^{-8} M\)
3. \(1.5 \times 10^{-8} M\)
4. \(1.0 \times 10^{-8} M\)

Given that the ionic product of $\mathrm{Ni}{\left(\mathrm{OH}\right)}_2$ is 2 x ${10}^{-\mathrm{15 }}$.
The solubility of $\mathrm{Ni}{\left(\mathrm{OH}\right)}_2$ in 0.1 M NaOH is ;

1. 2 x ${10}^{-8}$ M

2. 1 x ${10}^{-13}$ M

3. 1 x ${10}^8$ M 

4. 2 x ${10}^{-13}$ M

The molar solubility of ${\mathrm{CaF}}_2$ $({\mathrm{K}}_{\mathrm{sp}}=5.3$ $\mathrm{x}$ ${10}^{-11})$ in 0.1 M solution of NaF will be:

1. $5.3$ $\mathrm{x}$ ${10}^{-11}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$

2. $5.3$ $\mathrm{x}$ ${10}^{-8}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$

3. $5.3$ $\mathrm{x}$ ${10}^{-9}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$

4. $5.3$ $\mathrm{x}$ ${10}^{-10}$ $\mathrm{mol}$ ${\mathrm{L}}^{-1}$

The solubility of BaSO₄ in water is $2.42$ $\times$ ${10}^{-3}$ g/ litre at 298 K. The value of the solubility product will be: (Molar mass of BaSO₄ = 233 gmol^–1)

The concentration of Ag⁺ ions in a saturated solution of Ag₂C₂O₄ is 2.2 × 10^-4 mol L^-1.
The solubility product of Ag₂C₂O₄ is:

The solubility of AgCl (s) with solubility product 1.6 x 10^-10 in 0.1 M NaCl solution would be?

At room temperature, MY and NY₃, two nearly insoluble salts, have the same K_sp values of 6.2 x 10^-13. The true statement regarding MY and NY₃ is:

The K_sp of Ag₂CrO₄, AgCl, AgBr, and Agl are respectively, 1.1 x 10^-12, 1.8 x 10^-10, 5.0 x 10^-13, 8.3 x 10^-17. Which one of the following salts will precipitate last if ${\mathrm{AgNO}}_3$ solution is added to the solution containing equal moles of NaCl, NaBr, Nal, and Na₂CrO₄?

1. Agl

2. AgCl

3. AgBr

4. Ag₂CrO₄

pH of a saturated solution of Ba(OH)₂ is 12. The value of solubility product K_sp of Ba (OH)₂ is:
1. $3.3\times {10}^{-7}$
2. $5.0\times {10}^{-7}$
3. $4.0\times {10}^{-6}$
4. $5.0\times {10}^{-6}$

In qualitative analysis, the metals of Group I can be separated from other ions by precipitating them as chloride salts. A solution initially contains Ag⁺ and Pb²⁺ at a concentration of 0.10 M. Aqueous HCl is added to this solution until the Cl^– concentration is 0.10 M. What will the concentration of Ag⁺ and Pb²⁺ at equilibrium?

(K_sp for AgCl  = 1.8 × 10^-10)$$
(K_sp for PbCl₂ = 1.7 × 10^-5)