The ratio of F^- and HCOO^- in a mixture of 0.1 M HF $({\mathrm{K}}_{\mathrm{a}}=6.6\times {10}^{-4})$ and 0.2 M HCOOH $({\mathrm{K}}_{\mathrm{a}}=2\times {10}^{-4})is:$

1.	2 : 2.3	2.	1 : 3.3
3.	3.3 : 2	4.	3.3 : 1

A buffer solution is prepared by mixing 10 mL of 1.0 M acetic acid with 20 mL of 0.5 M sodium acetate which is then diluted
to 100 mL with distilled water. If the pK_a of ${\mathrm{CH}}_3\mathrm{COOH}$ is 4.76, the pH of the buffer solution prepared is -
1. 5.21

2. 4.76

3. 4.34

4. 5.35

What is the pH of a solution obtained by mixing equal volumes of two solutions, one with a pH of 3 and the other with a pH of 4?
[log 5.5=0.7404]

1. 3.26

2. 3.5

3. 4.0

4. 3.42

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

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