Solid AgNO₃ is added slowly to a buffer solution of pH=10 to precipitate AgOH. The [Ag⁺] concentration in the solution is $\left[{\mathrm{K}}_{\mathrm{sp}}\right(\mathrm{AgOH})={10}^{-10}]$

1. ${10}^{-4} \mathrm{M}$

2. ${10}^{-5} \mathrm{M}$

3. ${10}^{-6} \mathrm{M}$

4. ${10}^{-7} \mathrm{M}$

If pK_b for ${\mathrm{CN}}^{-}$ at $25°\mathrm{C}$ is 4.7, the pH of 0.5 M aqueous NaCN solution is:

1. 10

2. 11.5

3. 11

4. 12

A buffer solution is prepared by mixing 20 ml of 0.1 M CH₃COOH and 40 ml of 0.5 M CH₃COONa and then diluted by adding 100 ml of distilled water. The pH of the resulting buffer solution is (Given pKa CH₃COOH=4.76)

1. 5.76

2. 3.67

3. 2.48

4. 4.76

An acid-base indicator has ${\mathrm{K}}_{\mathrm{a}}=3.0\times {10}^{-5}$. The acid form of the indicator is red and the basic form is blue. The [H⁺] required to change the indicator from 75% red to 75% blue is

1. $8\times {10}^{-5} \mathrm{M}$

2. $9\times {10}^{-5} \mathrm{M}$

3. $1\times {10}^{-5} \mathrm{M}$

4. $3\times {10}^{-4} \mathrm{M}$

The self-ionization constant for pure formic acid, K = [\(H C O O H_{2}^{+} \left]\right. \left[\right. H C O O^{-} \left]\right.\) ] has been estimated as \(\left(10\right)^{- 4}\) at room temperature. The percentage of formic acid molecules in pure formic acid that are converted to formate ions is

\(\left(\right. Given : d_{HCOOH} = 1 . 22\ g / cc \left.\right)\)

1. 0.0185%

2. 0.0073%

3. 0.074%

4. 0.037%

AgOH is added to NaCl solution to form AgCl precipitate. After the precipitation, the pH of the solution is 8. The [Cl^-] is $({\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgCl}={10}^{-12}, {\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgOH}={10}^{-10})$

1. ${10}^{-6} \mathrm{M}$

2. ${10}^{-4} \mathrm{M}$

3. ${10}^{-8} \mathrm{M}$

4. ${10}^{-10} \mathrm{M}$

For the reaction 
${\left[\mathrm{Ag}{\left(\mathrm{CN}\right)}_2\right]}^{-}\left(\mathrm{aq}\right)\rightleftharpoons {\mathrm{Ag}}^{+}\left(\mathrm{aq}\right)+2{\mathrm{CN}}^{-}\left(\mathrm{aq}\right),$the equilibrium constant at $25°\mathrm{C}$ is $4.0\times {10}^{-19}.$ The silver ion concentration in a solution that was originally 0.10 M in KCN and 0.03 M in AgNO₃ will be:

1. $2.5\times {10}^{-18} \mathrm{M}$

2. $1.5\times {10}^{-18} \mathrm{M}$

3. $5.5\times {10}^{-18} \mathrm{M}$

4. $7.5\times {10}^{-18} \mathrm{M}$

Separate solutions of four sodium salts NaW, NaX, NaY and NaZ had pH 7.0, 9.0, 10.0 and 11.0 respectively. When each solution is 0.1 M, the strongest acid is

1. HW

2. HX

3. HY

4. HZ

Silver acetate is a slightly soluble salt of a weak acid $({\mathrm{K}}_{\mathrm{a}}=1.75\times {10}^{-5})$. At $25°\mathrm{C}$, 100 g of water dissolves 1.04 g of crystalline silver acetate. The density of saturated solution of silver acetate at $20°\mathrm{C}$ is 1.01 g/cc. The solubility product constant for silver acetate at $20°\mathrm{C\; will\; be:}$

1. $2.43\times {10}^{-3}$

2. $3.87\times {10}^{-3}$

3. $7.74\times {10}^{-5}$

4. $1.35\times {10}^{-5}$

0.1 M solution of three different sodium salts NaX, NaY and NaZ have pH values 7.0, 9.0 and 11.0 respectively. The correct order of dissociation constant values of these acids is

1. K_HX<K_HY<K_HZ

2. K_HX>K_HY>K_HZ

3. K_HX>K_HZ>K_HY

4. K_HX<K_HY<K_HZ