Chapter 7

When rain is accompanied by a thunder storm, the collected rain water will have a \(\mathrm{pH}\) value (a) depending on the amount of dust in air. (b) slightly lower than that of rain water without thunderstorm. (c) slightly higher than that when the thunder storm is not there. (d) uninfluenced by occurrence of thunderstorm.

For a sample of pure water, (a) pH increases and pOH decreases with increase in temperature. (b) pH decreases and pOH increases with increase in temperature. (c) both \(\mathrm{pH}\) and \(\mathrm{pOH}\) increases with increase in temperature. (d) both \(\mathrm{pH}\) and \(\mathrm{pOH}\) decrease with increase in temperature.

The \(\mathrm{pH}\) at which water is maximum dissociated at \(25^{\circ} \mathrm{C}\), is (a) 2 (b) 7 (c) 10 (d) 14

The degree of dissociation of water at \(25^{\circ} \mathrm{C}\) is \(1.8 \times 10^{-7} \%\) and density is \(1.0 \mathrm{~g} \mathrm{~cm}^{-3}\). The ionic constant for water is (a) \(1.0 \times 10^{-14}\) (b) \(2.0 \times 10^{-16}\) (c) \(1.0 \times 10^{-16}\) (d) \(1.0 \times 10^{-8}\)

The degree of dissociation of pure water at \(25^{\circ} \mathrm{C}\) is found to be \(1.8 \times 10^{-9}\). The dissociation constant, \(K_{\mathrm{d}}\) of water, at \(25^{\circ} \mathrm{C}\) is (a) \(10^{-14}\) (b) \(1.8 \times 10^{-16}\) (c) \(5.56 \times 10^{-13}\) (d) \(1.8 \times 10^{-14}\)

When \(20 \mathrm{ml}\) of \(0.2 \mathrm{M}-\mathrm{DCl}\) solution is mixed with \(80 \mathrm{ml}\) of \(0.1 \mathrm{M}-\mathrm{NaOD}\) solution, \(\mathrm{pD}\) of the resulting solution becomes \(13.6 .\) The ionic product of heavy water, \(\mathrm{D}_{2} \mathrm{O}\), is (a) \(10^{-15}\) (b) \(10^{-16}\) (c) \(4 \times 10^{-15}\) (d) \(4 \times 10^{-16}\)

What is the pH of a neutral solution at \(37^{\circ} \mathrm{C}\), where \(K_{w}\) equals \(2.5 \times 10^{-14} ?(\log 2=0.3)\) (a) \(7.0\) (b) \(13.6\) (c) \(6.8\) (d) \(6.6\)

The hydronium ion concentration in an aqueous solution of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is \(2.0 \times 10^{-4} \mathrm{M}\) at \(25^{\circ} \mathrm{C}\). The hydroxide ion concentration in the solution is (a) 0 (b) \(2.0 \times 10^{-4} \mathrm{M}\) (c) \(5 \times 10^{3} \mathrm{M}\) (d) \(5 \times 10^{-11} \mathrm{M}\)

The dissociation constant of a weak monoprotic acid is numerically equal to the dissociation constant of its conjugate base. What is \(\mathrm{pH}\) of \(0.1 \mathrm{M}\) solution of this acid? (a) \(7.0\) (b) \(6.0\) (c) \(8.0\) (d) \(4.0\)

The number of hydronium ions in \(1 \mathrm{ml}\) of an aqueous solution of \(\mathrm{pH} 12.0\) at \(25^{\circ} \mathrm{C}\) is (a) \(0.01\) (b) \(10^{-12}\) (c) \(6.02 \times 10^{8}\) (d) \(6.02 \times 10^{11}\)

The \(\mathrm{pH}\) of \(4.0 \times 10^{-4} \mathrm{M}-\mathrm{HNO}_{3}\) solution is \((\log 2=0.3)\) (a) \(4.6\) (b) \(3.4\) (c) \(3.6\) (d) \(4.0\)

A solution contains \(4.25 \mathrm{~g}\) ammonia per \(250.0 \mathrm{~m}\) lof solution. Electrical conductivity measurement at \(25^{\circ} \mathrm{C}\) show that \(0.40 \%\) of the ammonia has reacted with water. The pH of the solution is \((\log 2=0.3)\) (a) \(11.6\) (b) \(2.4\) (c) \(12.6\) (d) \(10.6\)

What mass of NaOH should be dissolved in sufficient water to get \(20 \mathrm{~m}^{3}\) of an aqueous solution of \(\mathrm{pH}, 7.3\), at \(25^{\circ} \mathrm{C}\) ? (a) \(0.16 \mathrm{~g}\) (b) \(1.6 \times 10^{-4} \mathrm{~g}\) (c) \(0.04 \mathrm{~g}\) (d) \(0.12 \mathrm{~g}\)

At \(25^{\circ} \mathrm{C}\), the dissociation constants of acid HA and base BOH in aqueous solution is same. The \(\mathrm{pH}\) of \(0.01 \mathrm{M}\) solution of HA is \(5.0\). The pH of \(0.1\) M solution of \(\mathrm{BOH}\) is (a) \(5.0\) (b) \(9.0\) (c) \(9.5\) (d) \(8.5\)

Following five solutions of \(\mathrm{KOH}\) were prepared as: first, \(0.1\) mole in \(1 \mathrm{~L} ;\) second. \(0.2\) mole in \(2 \mathrm{~L}\); third, \(0.3\) mole in \(3 \mathrm{~L}\); fourth, \(0.4\) mole in \(4 \mathrm{~L} ;\) fifth, \(0.5\) mole in \(5 \mathrm{~L}\). The \(\mathrm{pH}\) of resultant solution, when all these solutions are mixed, is (a) 2 (b) 1 (c) 13 (d) 7

An aqueous solution initially contains \(0.01 \mathrm{M}-\mathrm{RNH}_{2}\left(K_{\mathrm{b}}=2.0 \times 10^{-6}\right)\) and \(10^{-4} \mathrm{M}-\mathrm{NaOH}\). The final concentration of \(\mathrm{OH}^{-}\) in the solution is about (a) \(10^{-4} \mathrm{M}\) (b) \(2.0 \times 10^{-4} \mathrm{M}\) (c) \(3.0 \times 10^{-4} \mathrm{M}\) (d) \(1.414 \times 10^{-4} \mathrm{M}\)

What will be the effect of adding \(100 \mathrm{ml}\) of \(0.001 \mathrm{M}-\mathrm{HCl}\) solution to \(100 \mathrm{ml}\) of a solution having \(0.1 \mathrm{M}-\mathrm{HA}\) ? The acid dissociation constant of \(\mathrm{HA}\) is \(10^{-5}\). (a) The degree of dissociation of HA will decrease but the \(\mathrm{pH}\) of solution remains unchanged. (b) The degree of dissociation of \(\mathrm{HA}\) remains unchanged but the \(\mathrm{pH}\) of solution decreases. (c) Neither degree of dissociation nor pH of solution will change. (d) The degree of dissociation as well as pH of solution will decrease.

When \(0.05\) moles of the following acid are dissolved in \(1000 \mathrm{ml}\) of \(\mathrm{H}_{2} \mathrm{O}\), the \(\left[\mathrm{H}^{+}\right.\) will be greatest in (a) \(\mathrm{HNO}_{2} ; \mathrm{p} K_{\mathrm{a}}=3.0\) (b) \(\mathrm{HCOOH} ; \mathrm{p} K_{\mathrm{a}}=3.75\) (c) \(\mathrm{HCN} ; \mathrm{p} K_{\mathrm{a}}=9.4\) (d) \(\mathrm{CH}_{3} \mathrm{COOH} ; \mathrm{p} K_{\mathrm{a}}=4.75\)

Water in equilibrium with air contains \(4.4 \times 10^{-5} \% \mathrm{CO}_{2}\). The resulting carbonic acid, \(\mathrm{H}_{2} \mathrm{CO}_{3}\), gives the solution a hydronium ion concentration of \(2.0\) \(\times 10^{-6} \mathrm{M}\), about 20 times greater than that of pure water. What is the \(\mathrm{pH}\) of the solution at \(298 \mathrm{~K} ?(\log 4.4=0.64\) \(\log 2=0.3\) ) (a) \(5.36\) (b) \(5.70\) (c) \(8.30\) (d) \(5.64\)

The concentration of acetate ions in \(1 \mathrm{M}\) acetic acid \(\left(K_{\mathrm{a}}=2 \times 10^{-5}\right)\) solution containing \(0.1 \mathrm{M}-\mathrm{HCl}\) is (a) \(2 \times 10^{-1} \mathrm{M}\) (b) \(2 \times 10^{-3} \mathrm{M}\) (c) \(2 \times 10^{-4} \mathrm{M}\) (d) \(4.4 \times 10^{-3} \mathrm{M}\)c

What is the \(\mathrm{pH}\) of \(6.67 \times 10^{-3} \mathrm{M}\) aqueous solution of \(\mathrm{Al}(\mathrm{OH})_{3}\) if its first dissociation is \(100 \%\), second dissociation is \(50 \%\) and the third dissociation is negligible. (a) 2 (b) 12 (c) 11 (d) 3

The dissociation constant of acetic acid is \(0.000018\) and that for cyanoacetic acid is \(0.0036\) at \(298 \mathrm{~K}\). What would be the ratio of volumes of the two acid solutions, each containing equal moles of the acids, so that the solutions becomes isohydric? (a) \(1: 1\) (b) \(1: \sqrt{200}\) (c) \(1: 200\) (d) \(200: 1\)

Calculate \(\left[\mathrm{S}^{2}\right]\) in a solution originally having \(0.1 \mathrm{M}-\mathrm{HCl}\) and \(0.2 \mathrm{M}-\mathrm{H}_{2} \mathrm{~S}\). For \(\mathrm{H}_{2} \mathrm{~S}, K_{\mathrm{al}}=1.4 \times 10^{-7}\) and \(K_{\mathrm{a} 2}=1.0 \times 10^{-14}\). (a) \(0.1 \mathrm{M}\) (b) \(2.8 \times 10^{-20} \mathrm{M}\) (c) \(2.8 \times 10^{-22} \mathrm{M}\) (d) \(1.4 \times 10^{-20} \mathrm{M}\)

For a tribasic acid, \(\mathrm{H}_{3} \mathrm{~A}, K_{\mathrm{al}}=2 \times 10^{-5}\), \(K_{\mathrm{a} 2}=5 \times 10^{-9}\) and \(K_{\mathrm{a} 3}=4 \times 10^{-12}\). The value of \(\frac{\left[\mathrm{A}^{3-}\right]}{\left[\mathrm{H}_{3} \mathrm{~A}\right]}\) at equilibrium in an aqueous solution originally having \(0.2 \mathrm{M}-\mathrm{H}_{3} \mathrm{~A}\) is (a) \(5 \times 10^{-17}\) (b) \(5 \times 10^{-9}\) (c) \(1 \times 10^{-17}\) (d) \(2 \times 10^{-22}\)

To \(20 \mathrm{ml}\) of \(0.1 \mathrm{M}-\mathrm{NaOH}\) solution, \(3 \mathrm{ml}\) of \(1 \mathrm{M}\) acetic acid solution is added. Is the solution now neutral, acidic or alkaline? How much more of the acetic acid solution we add to produce a change of \(\mathrm{pH}=0.3\) unit? \(\left(\mathrm{p} K_{\mathrm{a}}\right.\) for \(\mathrm{CH}_{3} \mathrm{COOH}\) \(=4.74, \log 2=0.3\) ) (a) acidic, \(2 \mathrm{ml}\) (b) alkaline, \(1 \mathrm{ml}\) (c) acidic, \(1 \mathrm{ml}\) (d) neutral, \(2 \mathrm{ml}\)

A volume of \(18 \mathrm{ml}\) of mixture of acetic acid and sodium acetate required \(6 \mathrm{ml}\) of \(0.1 \mathrm{M}-\mathrm{NaOH}\) for neutralization of the acid and \(12 \mathrm{ml}\) of \(0.1 \mathrm{M}-\mathrm{HCl}\) reaction with salt separately. If \(\mathrm{p} K_{\mathrm{a}}\) of acetic acid is \(4.75\), what is the \(\mathrm{pH}\) of the mixture? \((\log 2=0.3)\) (a) \(5.05\) (b) \(4.45\) (c) \(4.15\) (d) \(5.35\)

To a solution of acetic acid, solid sodium acetate is gradually added. When ' \(\mathrm{x} \mathrm{g}\) ' of the salt has been added, the \(\mathrm{pH}\) has a certain value. When total 'y g' of the salt has been added, the \(\mathrm{pH}\) has been further raised by \(0.6\) units. What is the ratio of \(x: y ?(\log 3.98=0.6)\) (a) \(3.98: 1\) (b) \(1: 3.98\) (c) \(2: 3.98\) (d) \(3.98: 2\)

Two buffers, \(X\) and \(Y\) of \(p H 4.0\) and \(6.0\) respectively are prepared from acid HA and the salt NaA. Both the buffers are \(0.50 \mathrm{M}\) in HA. What would be the pH of the solution obtained by mixing equal volumes of the two buffers? \(K_{\mathrm{a}}\) of \(\mathrm{HA}=1.0 \times 10^{-5} \cdot(\log 5.05=0.7)\) (a) \(5.0\) (b) \(4.3\) (c) \(4.7\) (d) \(5.7\)

The dissociation constant of formic acid is \(0.00024\). The hydrogen ion concentration in \(0.002 \mathrm{M}-\mathrm{HCOOH}\) solution is nearly (a) \(6.93 \times 10^{-4} \mathrm{M}\) (b) \(4.8 \times 10^{-7} \mathrm{M}\) (c) \(5.8 \times 10^{-4} \mathrm{M}\) (d) \(1.4 \times 10^{-4} \mathrm{M}\)

The buffer capacity \((\beta)\) for a weak acid (A) \(-\) conjugate base (B) buffer is defined as the number of moles of strong acid or base needed to change the \(\mathrm{pH}\) of \(1 \mathrm{~L}\) of solution by \(1 \mathrm{pH}\) unit, where \(\beta=\frac{2.303\left(C_{\mathrm{A}}+C_{\mathrm{B}}\right) K_{\mathrm{a}}\left[\mathrm{H}^{+}\right]}{\left(\left[\mathrm{H}^{+}\right]+K_{\mathrm{a}}\right)^{2}} .\) Under what condition will a buffer best resist a change in \(\mathrm{pH}\) ? (a) \(\mathrm{pH}=3 \mathrm{p} \mathrm{Ka}\) (b) \(2 \mathrm{pH}=\mathrm{p} \mathrm{Ka}\) (c) \(\mathrm{pH}=\mathrm{p} \mathrm{Ka}\) (d) \(\mathrm{pH}=2 \mathrm{p} \mathrm{Ka}\)

Calculate pH of \(0.02 \mathrm{M}-\) HA solution. \(K_{\mathrm{a}}\) for \(\mathrm{HA}=2 \times 10^{-12} .(\log 2=0.3\) \(\log 3=0.48\) ) (a) \(6.65\) (b) \(6.70\) (c) \(6.85\) (d) \(6.52\)

A \(40.0 \mathrm{ml}\) solution of weak base, \(\mathrm{BOH}\) is titrated with \(0.1 \mathrm{~N}-\mathrm{HCl}\) solution. The \(\mathrm{pH}\) of the solution is found to be \(10.0\) and \(9.0\) after adding \(5.0 \mathrm{ml}\) and \(20.0 \mathrm{ml}\) of the acid, respectively. The dissociation constant of the base is \((\log 2=0.3)\) (a) \(2 \times 10^{-5}\) (b) \(1 \times 10^{-5}\) (c) \(4 \times 10^{-5}\) (d) \(5 \times 10^{-5}\)

How much water must added to \(300 \mathrm{ml}\) of \(0.2 \mathrm{M}\) solution of \(\mathrm{CH}_{3} \mathrm{COOH}\) for the degree of dissociation of the acid to double? \(K_{\mathrm{a}}\) for the acetic acid \(=1.8 \times 10^{-5}\). (a) \(1200 \mathrm{ml}\) (b) \(300 \mathrm{ml}\) (c) \(600 \mathrm{ml}\) (d) \(900 \mathrm{ml}\)

How many grams of \(\mathrm{NaOH}\) should be added in \(500 \mathrm{ml}\) of \(2 \mathrm{M}\) acetic acid solution to get a buffer solution of maximum buffer capacity? (a) \(20.0\) (b) \(10.0\) (c) \(40.0\) (d) \(30.0\)

A \(0.28 \mathrm{~g}\) sample of an unknown monoprotic organic acid is dissolved in water and titrated with a 0.1 M sodium hydroxide solution. After the addition of \(17.5 \mathrm{ml}\) of base, a pH of \(5.0\) is recorded. The equivalence point is reached when a total of \(35.0 \mathrm{ml}\) of \(\mathrm{NaOH}\) is added. The molar mass of the organic acid is (a) 160 (b) 80 (c) 40 (d) 120

What is the \(\mathrm{pH}\) of \(4 \times 10^{-3} \mathrm{M}-\mathrm{Y}(\mathrm{OH})_{2}\) solution assuming the first dissociation to be \(100 \%\) and second dissociation to be \(50 \%\), where \(Y\) represents a metal cation? \((\log 2=0.3, \log 3=0.48)\) (a) \(11.78\) (b) \(11.22\) (c) \(2.22\) (d) \(2.78\)

An aqueous solution is prepared by dissolving \(0.1\) mole \(\mathrm{H}_{2} \mathrm{CO}_{3}\) in sufficient water to get \(100 \mathrm{ml}\) solution at \(25^{\circ} \mathrm{C}\). For \(\mathrm{H}_{2} \mathrm{CO}_{3}, \quad K_{\mathrm{a} 1}=4.0 \times 10^{-6}\) and \(K_{\mathrm{a} 2}=5.0 \times 10^{-11} .\) The only incorrect equilibrium concentration is (a) \(\left[\mathrm{H}^{+}\right]=6.32 \times 10^{-4} \mathrm{M}\) (b) \(\left[\mathrm{HCO}_{3}\right]=2 \times 10^{-3} \mathrm{M}\) (c) \(\left[\mathrm{CO}_{3}^{2-}\right]=5 \times 10^{-11} \mathrm{M}\) (d) \(\left[\mathrm{OH}^{-}\right]=5 \times 10^{-12} \mathrm{M}\)

What is the aqueous ammonia concentration of a solution prepared by dissolving \(0.15\) mole of \(\mathrm{NH}_{4}^{+} \mathrm{CH}_{3} \mathrm{COO}^{-}\) in 1 L of water? Given: \(K_{\text {a }}\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) \(=1.8 \times 10^{-5} ; K_{\mathrm{b}}\left(\mathrm{NH}_{4} \mathrm{OH}\right)=1.8 \times 10^{-5}\) (a) \(8.3 \times 10^{-4} \mathrm{M}\) (b) \(0.15 \mathrm{M}\) (c) \(5.52 \times 10^{-3} \mathrm{M}\) (d) \(3.8 \times 10^{-4} \mathrm{M}\)

Ascorbic acid (vitamin \(\mathrm{C}\) ) is a diprotic acid, \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{6}\). What is the \(\mathrm{pH}\) of a \(0.10 \mathrm{M}\) solution? The acid ionization constants are \(K_{\mathrm{al}}=9.0 \times 10^{-5}\) and \(K_{\mathrm{a} 2}=1.6 \times 10^{-12} \cdot(\log 2=0.3, \log 3=0.48)\) (a) \(3.52\) (b) \(2.52\) (c) \(1.52\) (d) \(2.48\)

A volume of \(10 \mathrm{ml}\) of \(0.1 \mathrm{M}\) tribasic acid, \(\mathrm{H}_{3} \mathrm{~A}\) is titrated with \(0.1 \mathrm{M}-\mathrm{NaOH}\) solution. What is the ratio (approximate value) of \(\frac{\left[\mathrm{H}_{3} \mathrm{~A}\right]}{\left[\mathrm{A}^{3-}\right]}\) at the second equivalent point? Given: \(K_{1}=7.5 \times 10^{-4} ; K_{2}=10^{-8}\); \(K_{3}=10^{-12}\) (a) \(10^{-4}\) (b) \(10^{-3}\) (c) \(10^{-7}\) (d) \(10^{-6}\)

The dissociation constant of a weak acid \(\mathrm{HX}\) is, \(10^{-5}\). The buffer \(\mathrm{HX}+\mathrm{NaX}\) can be best used to maintain the \(\mathrm{pH}\) in the range (a) \(9-11\) (b) \(2-4\) (c) \(11-13\) (d) \(4-6\)

\(\begin{array}{ll}\text { The } \text { equilibrium } & \text { carbonate } \text { ion }\end{array}\) concentration after equal volumes of \(0.7 \mathrm{M}-\mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(0.7 \mathrm{M}-\mathrm{HCl}\) solutions are mixed, is \(\left(K_{\mathrm{al}}\right.\) and \(K_{\mathrm{a} 2}\) for \(\mathrm{H}_{2} \mathrm{CO}_{3}\) are \(4.9 \times 10^{-6}\) and \(4.0 \times 10^{-11}\), respectively) (a) \(0.7 \mathrm{M}\) (b) \(0.35 \mathrm{M}\) (c) \(0.002 \mathrm{M}\) (d) \(0.001 \mathrm{M}\)

Calcium Lactate is a salt of weak acid and represented as \(\mathrm{Ca}(\mathrm{Lac})_{2} .\) A saturated solution of \(\mathrm{Ca}(\mathrm{Lac})_{2}\) contains \(0.125 \mathrm{~mole}\) of salt in \(0.50 \mathrm{~L}\) solution. The \(\mathrm{pOH}\) of this is 5.60. Assuming complete dissociation of salt, calculate \(K_{\mathrm{a}}\) of lactate acid. \((\log 2.5=0.4)\) (a) \(1.25 \times 10^{-11}\) (b) \(8.0 \times 10^{-4}\) (c) \(3.2 \times 10^{-17}\) (d) \(4 \times 10^{-5}\)

The addition of sodium acetate to acetic acid solution will cause (a) increase in its \(\mathrm{pH}\) value (b) decrease in its \(\mathrm{pH}\) value (c) no change in \(\mathrm{pH}\) value (d) change in \(\mathrm{pH}\) which cannot be predicted

A \(0.1 \mathrm{M}\) acetic acid solution is titrated against \(0.1 \mathrm{M}-\mathrm{NaOH}\) solution. What would be the difference in \(\mathrm{pH}\) between \(1 / 4\) and \(3 / 4\) stages of neutralization of the acid? (a) \(2 \log (0.75)\) (b) \(2 \log (0.25)\) (c) \(\log 3\) (d) \(2 \log 3\)

When glycinium hydrochloride \(\left(\mathrm{NH}_{2} \mathrm{CH}_{2}\right.\) COOH.HCl) is titrated against \(\mathrm{NaOH}\), \(\mathrm{pH}\) at the first half equivalence point is \(2.40\) and the \(\mathrm{pH}\) at second half equivalence point is \(9.60\). The \(\mathrm{pH}\) at first equivalence point is (a) \(2.40\) (b) \(9.60\) (c) \(6.00\) (d) \(7.20\)

An amount of \(0.1\) mole of \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) \(\left(K_{\mathrm{b}}=5 \times 10^{-4}\right)\) is mixed with \(0.08\) mole of \(\mathrm{HCl}\) and diluted to one litre. What will be the \(\mathrm{H}^{+}\) concentration in the solution? (a) \(1.25 \times 10^{-4} \mathrm{M}\) (b) \(8 \times 10^{-11} \mathrm{M}\) (c) \(1.6 \times 10^{-11} \mathrm{M}\) (d) \(2 \times 10^{-3} \mathrm{M}\)

An acid base indicator which is a weak acid has a \(\mathrm{p} K_{\mathrm{a}}\) value \(=5.5\). At what concentration ratio of sodium acetate to acetic acid would the indicator show a colour half way between those of its acid and conjugate base forms? \(\mathrm{p} K_{\mathrm{a}}\) of acetic acid \(=4.75 .[\) Antilog \((0.75)=5.62\), Antilog \((0.79)=6.3\), Antilog \((0.69)=4.93]\) (a) \(4.93: 1\) (b) \(6.3: 1\) (c) \(5.62: 1\) (d) \(2.37: 1\)

An amount of \(0.15\) mole of pyridinium chloride has been added into \(500 \mathrm{ml}\) of 0.2 M pyridine solution. Calculate pH and hydroxyl ion concentration in the resulting solution assuming no change in volume. \(K_{\mathrm{b}}\) for pyridine \(=1.5 \times 10^{-9}\). \((\log 2=0.3, \log 0.3=0.48)\) (a) \(9.0\) (b) \(5.0\) (c) \(8.64\) (d) \(5.36\)

The correct increasing order of solubility of the following substances in \(\mathrm{g} / 100 \mathrm{ml}\) is \(\mathrm{PbSO}_{4}\left(K_{\mathrm{sp}}=2 \times 10^{-9}\right), \mathrm{ZnS}\left(K_{\mathrm{sp}}=1\right.\) \(\left.\times 10^{-22}\right)\), AgBr \(\left(K_{\text {sp }}=4 \times 10^{-13}\right), \mathrm{CuCO}_{3}\) \(\left(K_{\mathrm{sp}}=1 \times 10^{-8}\right) .\) (Atomic masses: \(\mathrm{Pb}=208\), \(\mathrm{Zn}=65, \mathrm{Ag}=108, \mathrm{Br}=80, \mathrm{Cu}=63)\) (a) \(\mathrm{PbSO}_{4}<\mathrm{ZnS}<\mathrm{AgBr}<\mathrm{CuCO}_{3}\) (b) \(\mathrm{PbSO}_{4}<\mathrm{CuCO}_{3}<\mathrm{AgBr}<\mathrm{ZnS}\) (c) \(\mathrm{ZnS}<\mathrm{AgBr}<\mathrm{CuCO}_{3}<\mathrm{PbSO}_{4}\) (d) \(\mathrm{ZnS}<\mathrm{AgBr}<\mathrm{PbSO}_{4}<\mathrm{CuCO}_{3}\)