Chapter 17, Problem 17.62SP

Interpretation Introduction

Interpretation:

The effect of salt addition on the pH of a weak base solution

Concept introduction:

• Buffer solution:Aqueous solution of a salt of strong acid and strong base e.g. sodium chloride (NaCl), and salt of weak acid and a weak base e.g. ammonium chloride ( $N{H}_{4}Cl$ ), both have a $pH$ of 7. But the addition of 1 cm³ of 0.1MHCl to 1 dm³ of the solutions alters the $pH$ to 4 in the case of former, while it hardly affects the $pH$ of the latter. The addition of an equivalent quantity of NaOH would like wise change the $pH$ of the NaCl solution from 7 to 10, but it would not appreciably alter that $N{H}_{4}Cl$ solution. The solution of $N{H}_{4}Cl$ thus has the property of resisting the change of $pH$
  
  when an acid or alkali is added to it and this property of solution is known as buffer action.
  
  In general, a buffer solution is one which is resistant to change of $pH$ upon the addition of a small amount of acid or base. The buffer solutions usually contain relatively larger amount of a weak base and salt of its conjugate acid. The $pOH$ of such basic the buffer solution is expressed by Henderson-Hasselbalch equation: $pOH=p{K}_b+\log \frac{[salt]}{[base]}$
• $pH$ and $pOH$: Since hydrogen-ion concentrations commonly met within solutions vary considerably over the range of ${10}^{-14}$ to 1M. Sorensen introduced a logarithmic scale for the sake of convenience, and gave the symbol $pH$. It is expressed as $pH=-\log C_{H^{+}}$. Similarly the concentration of the alkalinity of a solution is expressed by $pOH$ , which is expressed as $pOH=-\log C_{O{H}^{-}}$.

Given:

In 100mL 0.30M of $N{H}_3$ i.e. $N{H}_{4}OH$ ( $N{H}_3(aq)+H_{2}O(l)\to N{H}_{4}OH(aq)$ ) solution 4.0g of $N{H}_{4}N{O}_3$ is added. The volume of the solution remains constant. The base dissociation constant of $N{H}_{4}OH$ is $1.8\times {10}^{-5}$.

To determine:

The $pH$ of the $N{H}_3$ solution before and after the addition of $N{H}_{4}N{O}_3$.