Chapter 7

Separate solutions of NaW, NaX, NaY and NaZ, each of concentrations \(0.1 \mathrm{M}\), has \(\mathrm{pH} 7.0,9.0,10.0\) and \(11.0\) respectively, at \(25^{\circ} \mathrm{C}\). The strongest acid among these is (a) \(\mathrm{NaW}\) (b) NaX (c) NaY (d) \(\mathrm{NaZ}\)

If pH of \(0.001 \mathrm{M}\) potassium propionate solution be \(8.0\), then the dissociation constant of propionic acid will be (a) \(10^{-3}\) (b) \(10^{-2}\) (c) \(10^{-2.5}\) (d) \(10^{-5}\)

The correct order of increasing \(\left[\mathrm{OH}^{-}\right]\) in the following aqueous solution is (a) \(0.01 \mathrm{M}-\mathrm{NaHCO}_{3}<0.01 \mathrm{M}-\mathrm{NaCN}\) \(<0.01 \mathrm{M}-\mathrm{KCl}\) (b) \(0.01 \mathrm{M}-\mathrm{KCl}<0.01 \mathrm{M}-\mathrm{NaCN}\) \(<0.01 \mathrm{M}-\mathrm{NaHCO}_{3}\) (c) \(0.01 \mathrm{M}-\mathrm{KCl}<0.01 \mathrm{M}-\mathrm{NaHCO}_{3}\) \(<0.01 \mathrm{M}-\mathrm{NaCN}\) (d) \(0.01 \mathrm{M}-\mathrm{NaCN}<0.01 \mathrm{M}-\mathrm{KCl}\) \(<0.01 \mathrm{M}-\mathrm{NaHCO}_{3}\)

The solubility product of \(\mathrm{Co}(\mathrm{OH})_{3}\) is \(2.7 \times 10^{-43}\). The pH of saturated solution of \(\mathrm{Co}(\mathrm{OH})_{3}\) is about (a) \(7.0\) (b) \(11.0\) (c) \(3.0\) (d) \(3.48\)

In an attempted determination of the solubility product constant of \(\mathrm{Tl}_{2} \mathrm{~S}\), the solubility of this compound in pure \(\mathrm{CO}_{2}\) free water was determined as \(2.0 \times 10^{-6} \mathrm{M}\). Assume that the dissolved sulphide hydrolyses almost completely to \(\mathrm{HS}^{-}\) and that the further hydrolysis to \(\mathrm{H}_{2} \mathrm{~S}\) can be neglected, what is the computed \(K_{\mathrm{sp}}\) ? For \(\mathrm{H}_{2} \mathrm{~S}, K_{\mathrm{al}}=1.4 \times 10^{-7}, K_{\mathrm{a} 2}=1.0 \times 10^{-14}\) (a) \(6.4 \times 10^{-23}\) (b) \(1.6 \times 10^{-23}\) (c) \(3.2 \times 10^{-17}\) (d) \(3.2 \times 10^{-24}\)

After solid \(\mathrm{SrCO}_{3}\) was equilibrated with a buffer at \(\mathrm{pH} 8.6\), the solution was found to have \(\left[\mathrm{Sr}^{2+}\right]=2.0 \times 10^{-4} \mathrm{M}\), what is \(K_{\mathrm{sp}}\) of \(\mathrm{SrCO}_{3} ?\left(K_{\mathrm{a} 2}\right.\) for \(\mathrm{H}_{2} \mathrm{CO}_{3}=5.0 \times 10^{-11}\), \(\log 2=0.3,5.1 \times 0.196=1.0)\) (a) \(4.0 \times 10^{-8}\) (b) \(8.0 \times 10^{-8}\) (c) \(8.0 \times 10^{-10}\) (d) \(3.38 \times 10^{-8}\)

What is the solubility of \(\mathrm{MnS}\) in pure water, assuming hydrolysis of \(\mathrm{S}^{2-}\) ions? \(K_{\mathrm{sp}}\) of \(\mathrm{MnS}=2.5 \times 10^{-10}, K_{\mathrm{al}}=1 \times 10^{-7}\) and \(K_{\mathrm{a} 2}=1 \times 10^{-14}\) for \(\mathrm{H}_{2} \mathrm{~S} .\left(0.63^{3}=0.25\right)\) (a) \(6.3 \times 10^{-4} \mathrm{M}\) (b) \(2.5 \times 10^{-4} \mathrm{M}\) (c) \(6.3 \times 10^{-3} \mathrm{M}\) (d) \(1.58 \times 10^{-5} \mathrm{M}\)

The pH of \(0.1\) M solution of the following compounds increases in the order (a) \(\mathrm{NaCl}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCN}<\mathrm{HCl}\) (b) \(\mathrm{HCl}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCl}<\mathrm{NaCN}\) (c) \(\mathrm{NaCN}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCl}<\mathrm{HCl}\) (d) \(\mathrm{HCl}<\mathrm{NaCl}<\mathrm{NaCN}<\mathrm{NH}_{4} \mathrm{Cl}\)

An amount of \(0.10\) moles of \(\mathrm{AgCl}(\mathrm{s})\) is added to one litre of water. Next, the crystals of NaBr are added until \(75 \%\) of the \(\mathrm{AgCl}\) is converted to \(\mathrm{AgBr}(\mathrm{s})\), the less soluble silver halide. What is \(\mathrm{Br}^{-}\) at this point? \(K_{\mathrm{sp}}\) of \(\mathrm{AgCl}=2 \times 10^{-10}\) and \(K_{\mathrm{sp}}\) of \(\mathrm{AgBr}=4 \times 10^{-13}\) (a) \(0.075 \mathrm{M}\) (b) \(0.025 \mathrm{M}\) (c) \(1.5 \times 10^{-4} \mathrm{M}\) (d) \(0.027 \mathrm{M}\)

The pH value of \(0.1 \mathrm{M}\) solutions of \(\begin{array}{llll}\mathrm{CH}_{3} \mathrm{COONa} & \text { (I), } & \mathrm{CH}_{3} \mathrm{COOH} & \text { (II), }\end{array}\) \(\mathrm{CH}_{3} \mathrm{COONH}_{4}\) (III), \(\mathrm{NaOH}\) (IV) and \(\mathrm{HCl}(\mathrm{V})\) is in the order (a) \(\mathrm{I}<\mathrm{II}<\mathrm{III}<\mathrm{IV}<\mathrm{V}\) (b) \(\mathrm{V}<\mathrm{IV}<\mathrm{III}<\mathrm{II}<\mathrm{I}\) (c) \(\mathrm{V}<\mathrm{II}<\mathrm{III}<\mathrm{I}<\mathrm{IV}\) (d) \(\mathrm{V}<\mathrm{II}<\mathrm{I}<\mathrm{III}<\mathrm{IV}\)

An amount of \(0.01\) moles of solid \(\mathrm{AgCN}\) is rendered soluble in 11 by adding just sufficient excess cyanide ion to form \(\mathrm{Ag}(\mathrm{CN})_{2}^{-}\) and the concentration of free cyanide ion is \(2.5 \times 10^{-7} \mathrm{M}\). Determine \(\left[\mathrm{Ag}^{+}\right]\) in the solution neglecting hydrolysis of cyanide ion. \(K_{\text {diss }}\) for \(\mathrm{Ag}(\mathrm{CN})_{2}^{-}\) \(=1.0 \times 10^{-20}\) (a) \(6.25 \times 10^{-9} \mathrm{M}\) (b) \(1.6 \times 10^{-9} \mathrm{M}\) (c) \(1.6 \times 10^{-7} \mathrm{M}\) (d) \(6.25 \times 10^{-7} \mathrm{M}\)

Solid \(\mathrm{BaF}_{2}\) is added to a solution containing \(0.1\) mole of \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) solution (1 L) until equilibrium is reached. If the \(K_{\text {sp }}\) of \(\mathrm{BaF}_{2}\) and \(\mathrm{BaC}_{2} \mathrm{O}_{4}\) is \(10^{-6} \mathrm{~mol}^{3} \mathrm{~L}^{-3}\) and \(10^{-10} \mathrm{~mol}^{2} \mathrm{~L}^{-2}\), respectively, find the equilibrium concentration of \(\mathrm{Ba}^{2+}\) in the solution. Assume addition of \(\mathrm{BaF}_{2}\) does not cause any change in volume. (a) \(0.2 \mathrm{M}\) (b) \(4 \times 10^{-6} \mathrm{M}\) (c) \(2.5 \times 10^{-5} \mathrm{M}\) (d) \(2.5 \times 10^{-6} \mathrm{M}\)

What is the pH of a \(0.50 \mathrm{M}\) aqueous \(\mathrm{NaCN}\) solution? \(\mathrm{p} K_{\mathrm{b}}\) of \(\mathrm{CN}^{-}\) is \(4.70 .\) \((\log 2=0.3)\) (a) \(3.0\) (b) \(11.0\) (c) \(4.7\) (d) \(9.3\)

The \(\mathrm{pH}\) at the equivalence point when a solution of \(0.01 \mathrm{M}-\mathrm{CH}_{3} \mathrm{COOH}\) is titrated with a solution of \(0.01 \mathrm{M}-\mathrm{NaOH}\), is \(\left(\mathrm{p} K_{\mathrm{a}}\right.\) of \(\mathrm{CH}_{3} \mathrm{COOH}=4.7, \log 5=0.7\) ) (a) \(10.5\) (b) \(3.5\) (c) \(10.35\) (d) \(3.65\)

The concentration of \(\mathrm{CH}_{3} \mathrm{COO}^{-}\) ion in a solution prepared by adding \(0.1\) mole of \(\mathrm{CH}_{3} \mathrm{COOAg}(\mathrm{s})\) in \(1 \mathrm{~L}\) of \(0.1 \mathrm{M}-\mathrm{HCl}\) solution is [Given: \(K_{\mathrm{a}}\left(\mathrm{CH}_{3} \mathrm{COOH}\right)=10^{-5}\); \(\left.K_{\mathrm{sp}}(\mathrm{AgCl})=10^{-10} ; K_{\mathrm{sp}}\left(\mathrm{CH}_{3} \mathrm{COOAg}\right)=10^{-8}\right]\) (a) \(10^{-3} \mathrm{M}\) (b) \(10^{-2} \mathrm{M}\) (c) \(10^{-1} \mathrm{M}\) (d) \(1 \mathrm{M}\)

For the indicator thymol blue, the value of \(\mathrm{pH}\) is \(2.0 \mathrm{when}\) half of the indicator is present in the unionized form. The percentage of the indicator in the unionized form in a solution of \(4.0\) \(\times 10^{-3} \mathrm{M}\) hydrogen ion concentration is (a) \(40 \%\) (b) \(28.6 \%\) (c) \(71.4 \%\) (d) \(60 \%\)

Among different types of salts have nearly same solubility product constant, \(K_{\mathrm{sn}}\) but much smaller than one, the most soluble salt is that which (a) produces maximum number of ions (b) produces minimum number of ions (c) produces high charge on ions (d) produces low charges on ions

A certain sample of rainwater gives a yellow colour with methyl red [pH range \(4.2(\) red \()-6.2(\) yellow \()\) ]and a yellow colour with phenol red [pH range \(6.4\) (yellow) \(-8.0\) (red)]. What is the approximate \(\mathrm{pH}\) of the water? Is the rainwater acidic, neutral, or basic? (a) \(6.3\), acidic (b) \(6.1\), acidic (c) \(6.5\), acidic (d) \(6.3\), basic

An acid type indicator, HIn differs in colour from its conjugate base \(\left(\mathrm{In}^{-}\right)\). The human eye is sensitive to colour differences only when the ratio \([\operatorname{In}] /[\mathrm{HIn}]\) is greater than 10 or smaller than \(0.1\). What should be the minimum change in the \(\mathrm{pH}\) of the solution to observe a complete colour change \(\left(K_{\mathrm{a}}=1.0 \times 10^{-5}\right)\) ? (a) \(0.0\) (b) \(1.0\) (c) \(2.0\) (d) \(5.0\)

At what minimum pH will \(10^{-3}\) M \(-\mathrm{Al}(\mathrm{OH})_{3}\) go into solution \((V=1 \mathrm{~L})\) as \(\mathrm{Al}(\mathrm{OH})_{4}^{-}\) and at what maximum \(\mathrm{pH}\), it will dissolved as \(\mathrm{Al}^{3+}\) ? Given: \(\log 2=0.3\) \(\mathrm{Al}(\mathrm{OH})_{4}^{-} \rightleftharpoons \mathrm{Al}^{3+}+4 \mathrm{OH}^{-} ; K_{\mathrm{eq}}=1.6 \times 10^{-34}\) \(\mathrm{Al}(\mathrm{OH})_{3} \rightleftharpoons \mathrm{Al}^{3+}+3 \mathrm{OH}^{-} ; K_{\mathrm{eq}}=8.0 \times 10^{-33}\) (a) \(9.3,9.7\) (b) \(9.7,9.3\) (c) \(4.3,9.3\) (d) \(4.7,9.3\)

The solubility of sparingly soluble salt \(\mathrm{A}_{3} \mathrm{~B}_{2}\) (molar mass \(=\) 'M' \(\mathrm{g} / \mathrm{mol}\) ) in water is 'x' g/L. The ratio of molar concentration of \(\mathrm{B}^{3}\) to the solubility product of the salt is (a) \(\frac{108 x^{5}}{M^{5}}\) (b) \(\frac{x^{4}}{108 M^{4}}\) (c) \(\frac{x^{4}}{54 M^{4}}\) (d) \(\frac{x^{3}}{27 M^{3}}\)

The solubility product of \(\mathrm{AgCl}\) is \(1.0 \times 10^{-10}\). The equilibrium constant of the reaction \(\mathrm{AgCl}(\mathrm{s})+\mathrm{Br}^{-} \rightleftharpoons \mathrm{AgBr}(\mathrm{s})+\mathrm{Cl}\) is 200 and that of the reaction \(2 \mathrm{AgBr}(\mathrm{s})+\mathrm{S}^{2-} \rightleftharpoons \mathrm{Ag}_{2} \mathrm{~S}(\mathrm{~s})+2 \mathrm{Br}\) is \(1.6 \times 10^{24} .\) What is the \(K_{\mathrm{sp}}\) of \(\mathrm{Ag}_{2} \mathrm{~S}\) ? (a) \(3.2 \times 10^{16}\) (b) \(1.56 \times 10^{-49}\) (c) \(3.95 \times 10^{-25}\) (d) \(3.13 \times 10^{-17}\)

A solution contains \(0.1 \mathrm{M}-\mathrm{Mg}^{2+}\) and \(\begin{array}{lll}0.1 & \mathrm{M} & -\mathrm{Sr}^{2+} & \text { The concentration of }\end{array}\) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) in solution is adjusted to \(0.05 \mathrm{M}\). Determine the \(\mathrm{pH}\) range which would permit the precipitation of \(\mathrm{SrCO}_{3}\) without any precipitation of \(\mathrm{MgCO}_{3} . \mathrm{H}^{+}\) ion concentration is controlled by external factors. Given: \(K_{\mathrm{sp}}\left(\mathrm{MgCO}_{3}\right)=4 \times 10^{-8} \mathrm{M}^{2}\); \(K_{\mathrm{sp}}\left(\mathrm{SrCO}_{3}\right)=9 \times 10^{-10} \mathrm{M}^{2} ; K_{\mathrm{a}, \text { overall }}\left(\mathrm{H}_{2} \mathrm{CO}_{3}\right)\) \(=5 \times 10^{-17} ; \log 2=0.3 ; \log 3=0.48 .\) (a) \(4.78\) to \(5.6\) (b) \(4.6\) to \(5.78\) (c) \(5.78\) to \(6.4\) (d) \(5.22\) to \(5.4\)

A volume of \(500 \mathrm{ml}\) of \(0.01 \mathrm{M}-\mathrm{AgNO}_{3}\) solution, \(250 \mathrm{ml}\) of \(0.02 \mathrm{M}-\mathrm{NaCl}\) solution and \(250 \mathrm{ml}\) of \(0.02 \mathrm{M}-\mathrm{NaBr}\) solution are mixed. The final concentration of bromide ion in the solution is \(\left(K_{\mathrm{sp}}\right.\) of \(\mathrm{AgCl}\) and \(\mathrm{AgBr}\) are \(10^{-10}\) and \(5 \times 10^{-13}\) respectively.) (a) \(0.01 \mathrm{M}\) (b) \(0.02 \mathrm{M}\) (c) \(0.005 \mathrm{M}\) (d) \(2.5 \times 10^{-5} \mathrm{M}\)

How many times solubility of \(\mathrm{CaF}_{2}\) is decreased in \(4 \times 10^{-3} \mathrm{M}-\mathrm{KF}(\mathrm{aq})\) solution as compared to pure water at \(25^{\circ} \mathrm{C}\). Given: \(K_{\text {sp }}\left(\mathrm{CaF}_{2}\right)=3.2 \times 10^{-11}\) (a) 50 (b) 100 (c) 500 (d) 1000

For a sparingly soluble salt \(\mathrm{A}_{\mathrm{p}} \mathrm{B}_{\mathrm{q}}\), the relationship of its solubility product \(\left(L_{\mathrm{s}}\right)\) with its solubility \((S)\) is (a) \(L_{\mathrm{s}}=S^{p+q} \cdot p^{p} \cdot q^{q}\) (b) \(L_{\mathrm{s}}=S^{p+q} \cdot p^{q} \cdot q^{p}\) (c) \(L_{\mathrm{S}}=S^{p q} \cdot p^{p} \cdot q^{q}\) (d) \(L_{\mathrm{s}}=S^{p q} \cdot(p q)^{p+q}\)

Solubility product constant \(\left(K_{\mathrm{sp}}\right)\) of salts of types \(\mathrm{MX}, \mathrm{MX}_{2}\) and \(\mathrm{M}_{3} \mathrm{X}\) at temperature, \(T\) are \(4.0 \times 10^{-8}, 3.2 \times 10^{-14}\) and \(2.7 \times 10^{-15}\), respectively. Solubilities (in \(\mathrm{M}\) ) of the salts at temperature, \(T\), are in the order (a) \(\mathrm{MX}>\mathrm{MX}_{2}>\mathrm{M}_{3} \mathrm{X}\) (b) \(\mathrm{M}_{3} \mathrm{X}>\mathrm{MX}_{2}>\mathrm{MX}\) (c) \(\mathrm{MX}_{2}>\mathrm{M}_{3} \mathrm{X}>\mathrm{MX}\) (d) \(\mathrm{MX}>\mathrm{M}_{3} \mathrm{X}>\mathrm{MX}_{2}\)

The solubility of \(\mathrm{AgCl}\) in water is \(0.001435 \mathrm{~g}\) per litre at \(15^{\circ} \mathrm{C}\). The solubility product of \(\mathrm{AgCl}\) is \((\mathrm{Ag}=108, \mathrm{Cl}=35.3)\) (a) \(10^{-5}\) (b) \(10^{-10}\) (c) \(2 \times 10^{-10}\) (d) \(10^{-9}\)

The solubility of \(\mathrm{Li}_{3} \mathrm{Na}_{3}\left(\mathrm{AlF}_{6}\right)_{2}\) is \(0.0744 \mathrm{~g}\) per \(100 \mathrm{ml}\) at \(298 \mathrm{~K}\). Calculate the solubility product of the salt. (Atomic masses: \(\mathrm{Li}=7, \mathrm{Na}=23, \mathrm{Al}=27, \mathrm{~F}=19)\) (a) \(2.56 \times 10^{-22}\) (b) \(2 \times 10^{-3}\) (c) \(7.46 \times 10^{-19}\) (d) \(3.46 \times 10^{-12}\)

The solubility product of \(\mathrm{CaF}_{2}\) is \(1.08\) \(\times 10^{-10}\). What mass of \(\mathrm{CaF}_{2}\) will dissolve in \(500 \mathrm{ml}\) water in order to make a saturated solution? \((\mathrm{Ca}=40, \mathrm{~F}=19)\) (a) \(3 \times 10^{-4} \mathrm{~g}\) (b) \(1.17 \times 10^{-2} \mathrm{~g}\) (c) \(1.17 \mathrm{mg}\) (d) \(3 \times 10^{-3} \mathrm{~g}\)

The solubility product of \(\mathrm{Mg}(\mathrm{OH})_{2}\) is \(9.0 \times 10^{-12}\). The \(\mathrm{pH}\) of an aqueous saturated solution of \(\mathrm{Mg}(\mathrm{OH})_{2}\) is \((\log 1.8=0.26, \log 3=0.48)\) (a) \(3.58\) (b) \(10.42\) (c) \(3.88\) (d) \(6.76\)

The solubility of \(\mathrm{AgCN}\) in a buffer solution of \(\mathrm{pH}=3.0\) is \(\left(K_{\mathrm{sp}}\right.\) of \(\mathrm{AgCN}\) \(=1.2 \times 10^{-16} ; K_{\mathrm{a}}\) of \(\left.\mathrm{HCN}=4.8 \times 10^{-10}\right)\) (a) \(1.58 \times 10^{-5} \mathrm{M}\) (b) \(2.0 \times 10^{-5} \mathrm{M}\) (c) \(1.58 \times 10^{-4} \mathrm{M}\) (d) \(2.5 \times 10^{-9} \mathrm{M}\)

\(\begin{array}{llll}\text { Solubility } & \text { products of } & \mathrm{Mg}(\mathrm{OH})_{2} \text { , }\end{array}\) \(\mathrm{Cd}(\mathrm{OH})_{2}, \mathrm{Al}(\mathrm{OH})_{3}\) and \(\mathrm{Zn}(\mathrm{OH})_{2}\) are \(4 \times 10^{-11}, 8 \times 10^{-6}, 8.5 \times 10^{-23}\) and \(1.8 \times 10^{-14}\), respectively. The cation that will precipitate first as hydroxide, on adding limited quantity of \(\mathrm{NH}_{4} \mathrm{OH}\) in a solution containing equimolar amount of metal cations, is (a) \(\mathrm{Al}^{3+}\) (b) \(\mathrm{Zn}^{2+}\) (c) \(\mathrm{Mg}^{2+}\) (d) \(\mathrm{Cd}^{2+}\)

Silver ions are slowly added in a solution with \(\left[\mathrm{Br}^{-}\right]=\left[\mathrm{Cl}^{-}\right]=\left[\mathrm{CO}_{3}^{2-}\right]=\left[\mathrm{AsO}_{4}^{3-}\right]\) \(=0.1 \mathrm{M}\). Which compound will precipitate first? (a) \(\operatorname{AgBr}\left(K_{\mathrm{sp}}=5 \times 10^{-13}\right)\) (b) \(\mathrm{AgCl}\left(K_{\mathrm{sp}}=1.8 \times 10^{-10}\right)\) (c) \(\mathrm{Ag}_{2} \mathrm{CO}_{3}\left(K_{\mathrm{sp}}=8.1 \times 10^{-12}\right)\) (d) \(\mathrm{Ag}_{3} \mathrm{PO}_{4}\left(K_{\mathrm{sp}}=1 \times 10^{-22}\right)\)

A \(0.1\) mole of \(\mathrm{AgNO}_{3}\) is dissolved in \(1 \mathrm{~L}\) of \(1 \mathrm{M}-\mathrm{NH}_{3} .\) If \(0.01\) mole of \(\mathrm{NaCl}\) is added to this solution, will \(\mathrm{AgCl}(\mathrm{s})\) precipitate? \(K_{\mathrm{sp}}\) for \(\mathrm{AgCl}=1.8 \times 10^{-10}\) and \(K_{\text {stab }}\) for \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}^{+}=1.6 \times 10^{7} .\) (a) Yes (b) No (c) Addition of \(\mathrm{NaCl}\) in any amount can never result precipitation. (d) Addition of even smaller amount of \(\mathrm{NaCl}\) may result precipitation.

The minimum mass of NaBr which should be added in \(200 \mathrm{ml}\) of \(0.0004 \mathrm{M}-\mathrm{AgNO}_{3}\) solution just to start the precipitation of AgBr. \(K_{\mathrm{sp}}\) of \(\mathrm{AgBr}=4 \times 10^{-13} \cdot(\mathrm{Br}=80)\) (a) \(1.0 \times 10^{-9} \mathrm{~g}\) (b) \(2 \times 10^{-10} \mathrm{~g}\) (c) \(2.06 \times 10^{-8} \mathrm{~g}\) (d) \(1.03 \times 10^{-7} \mathrm{~g}\)

A sample of hard water contains \(0.005\) mole of \(\mathrm{CaCl}_{2}\) per litre. What is the minimum concentration of \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) which must be exceeded for removing \(\mathrm{Ca}^{2+}\) ions from this water sample? The solubility product of \(\mathrm{CaSO}_{4}\) is \(2.4 \times 10^{-5}\). (a) \(4.8 \times 10^{-3} \mathrm{M}\) (b) \(1.2 \times 10^{-3} \mathrm{M}\) (c) \(0.0144 \mathrm{M}\) (d) \(2.4 \times 10^{-3} \mathrm{M}\)

A solution contains a mixture of \(\mathrm{Ag}^{+}\) \((0.10 \mathrm{M})\) and \(\mathrm{Hg}_{2}^{2+}(0.10 \mathrm{M})\), which are to be separated by selective precipitation. Calculate the maximum concentration of iodide ion at which one of them gets precipitated almost completely. What per cent of that metal ion is precipitated, before the start of precipitation of second metal ion? \(K_{\mathrm{sp}}(\mathrm{AgI})=8.5 \times 10^{-17}\) and \(K_{\mathrm{sp}}\left(\mathrm{Hg}_{2} \mathrm{I}_{2}\right)=2.5 \times 10^{-26}\) (a) \(5 \times 10^{-13} \mathrm{M}, 99.83 \%\) (b) \(8.5 \times 10^{-16} \mathrm{M}, 99.83 \%\) (c) \(2.5 \times 10^{-25} \mathrm{M}, 100 \%\) (d) \(5 \times 10^{-13} \mathrm{M}, 98.3 \%\)

Small amount of freshly precipitated \(\begin{array}{llll}\text { magnesium } & \text { hydroxides } & \text { are } & \text { stirred }\end{array}\) vigorously in a buffer solution containing \(0.25 \mathrm{M}\) of \(\mathrm{NH}_{4} \mathrm{Cl}\) and \(0.05 \mathrm{M}\) of \(\mathrm{NH}_{4} \mathrm{OH}\). \(\left[\mathrm{Mg}^{2}\right]\) in the resulting solution is \(\left(K_{\mathrm{b}}\right.\) for \(\mathrm{NH}_{4} \mathrm{OH}=2.0 \times 10^{-5}\) and \(K_{\mathrm{sp}}\) of \(\mathrm{Mg}(\mathrm{OH})_{2}\) \(\left.=8.0 \times 10^{-12}\right)\) (a) \(4 \times 10^{-6} \mathrm{M}\) (b) \(2 \times 10^{-6} \mathrm{M}\) (c) \(0.5 \mathrm{M}\) (d) \(2.0 \mathrm{M}\)