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 of CH₃COOH=4.76)

An acid-base indicator has ${\mathrm{K}}_{\mathrm{a}}=3.0\times {10}^{-5}$. The acidic 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. \(8 \times 10^{- 5} M \)

2. \(9 \times 10^{- 5}~ \text M\)

3. \(1 \times 10^{- 5}~ \text M\)

4. \(3 \times 10 ^{- 4}~ \text M\)

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:
\([K_{sp} ~(AgCl)=10^{-12}, K_{sp}~ (AgOH) =10^{-10} ]\)
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. $2.5\times {10}^{-18}<mspace\; width="1em"></mspace>\mathrm{M}$

2. $1.5\times {10}^{-18}<mspace\; width="1em"></mspace>\mathrm{M}$

3. $5.5\times {10}^{-18}<mspace\; width="1em"></mspace>\mathrm{M}$

4. $7.5\times {10}^{-18}<mspace\; width="1em"></mspace>\mathrm{M}$

Identify the strongest acid among the corresponding acids of NaW, NaX, NaY, and NaZ, given that their 0.1 M solutions have pH values of 7.0, 9.0, 10.0, and 11.0, respectively.

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

When equal volume of the following solutions are mixed, which of the following gives maximum precipitate?
\((K_{sp}~ \text {of} ~AgCl = 10^{-12})\) $$

1. \(10^{-4} \mathrm{M}~ \mathrm{Ag}^{+}~\text {and} ~10^{-4} ~\mathrm{M}~ \mathrm{Cl}^{-}\)
2. \(10^{-3} \mathrm{M} ~\mathrm{Ag}^{+} \text {and }10^{-3} \mathrm{M} ~\mathrm{Cl}^{-}\)
3. \(10^{-5} \mathrm{M}~ \mathrm{Ag}^{+}~\text {and }10^{-5} \mathrm{M} ~\mathrm{Cl}^{-}\)
4. \(10^{-6} \mathrm{M} ~\mathrm{Ag}^{+}~\text {and }~10^{-6} \mathrm{M} ~\mathrm{Cl}^{-}\)