20 mL of a solution of acetic acid required 28.4 mL of 0.1 M NaOH for its neutralization. A solution (X) was prepared by mixing 20 mL of the above acetic acid and 14.2 mL of 0.1 M NaOH solution. What is the pH of the solution $(\mathrm{X})$ ? $\left(\mathrm{pK}_{\mathrm{a}}\right.$ value of acetic acid is 4.75).

The first and second ionization constants of a weak dibasic acid H₂A are $8.1 \times 10^{-8}$ and $1.0 \times 10^{-13}$ respectively. 0.1 mol of H₂A was dissolved in 1 L of 0.1 M HCl solution. The concentration of HA^- in the resultant solution is:

At 25°C, 20.0 mL of 0.2 M weak monoprotic acid HX is titrated against 0.2 M NaOH. The pH of the solution (a) at the start of the titration (when NaOH has not been added) and (b) when 10 mL of NaOH is added respectively, are :

Given :
$K_a = 5 \times 10^{-4}$
$\mathrm{p}K_a = 3.3$
$\alpha \ll 1$

The solubility product constants of $\mathrm{Ag_2CrO_4}$ and $\mathrm{AgBr}$ are $32x$ and $4y$ respectively at 298 K.

The value of $$\left( \frac{\text{molarity of } \mathrm{Ag_2CrO_4}}{\text{molarity of } \mathrm{AgBr}} \right)$$ can be expressed as :