Buffer Solutions

The EDTA titrations are generally carried out at a $\mathrm{pH}$ of $10$ . Why is it necessary to buffer the $\mathrm{pH}$ at $10$ ?

An acid-base indicator which is a weak acid has a ${\mathrm{pK}}_{\mathrm{a}}$ 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? ${\mathrm{pK}}_{\mathrm{a}}$ of acetic acid $=4.75$

A chemist needs a buffer with $\mathrm{pH}$ $4.35.$ How many $\mathrm{mL}$ of pure acetic acid (density $=1.049 \mathrm{g}/\mathrm{mL}$ ) must be added to $465 \mathrm{mL}$ of  $0.0941 \mathrm{M} \mathrm{NaOH}$ solution to obtain such a buffer? $\left(\mathrm{pKa} {\mathrm{CH}}_3\mathrm{COOH} = 4.74\right)$

Give answer after multiplying with 10 and rounding off to the nearest integer value.

A quantity of $0.25 \mathrm{M} \mathrm{NaOH}$ is added to a solution containing $0.15$mole of acetic acid. The final volume of the solution is $375 \mathrm{mL}$ and the $\mathrm{pH}$ of the solution is $4.45$. $\left({\mathrm{pK}}_{\mathrm{a}} {\mathrm{CH}}_3\mathrm{COOH} = 4.47, {10}^{-0.02}=0.948\right)$

What was the original concentration of the acetic acid?

Round off your to the nearest integer.

A quantity of $0.25 \mathrm{M} \mathrm{NaOH}$ is added to a solution containing$0.15$ mole of acetic acid. The final volume of the solution is $375 \mathrm{mL}$ and the $\mathrm{pH}$ of the solution is$4.45$. $\left({\mathrm{pK}}_{\mathrm{a}} {\mathrm{CH}}_3\mathrm{COOH} = 4.47\right)$

How many $\mathrm{mL}$ of $\mathrm{NaOH}$ was added to the original solution. Answer to the nearest integer.