Chapter 10.10, Problem 10.26P

Interpretation Introduction

Interpretation:

The $\text{pH}$ for given buffer solution with $0.050MN{H}_{4}Cland0.0800MN{H}_3$ should be determined.

Concept Introduction:

$\text{pH}$: $\text{pH}$ is the logarithm of the reciprocal of the concentration of ${\text{H}}_{\text{3}}{\text{O}}^{\text{+}}$ in a solution. The pH of a solution is a measure of concentration of hydrogen ion in a solution. Lower $\text{pH}$ due to more hydrogen ions and higher $\text{pH}$ due to less concentration of hydrogen ions.

$\text{pH}\text{=}{\text{pK}}_{\text{a}}\text{+}\text{log}\frac{\text{[conjugate base]}}{\text{[acid]}}$ is Henderson-Hasselbalch equation.

Henderson-Hasselbalch equation

$\text{pKa}\text{=}\text{pH}\text{+}\text{log}\frac{\left[\text{HA}\right]}{\left[{\text{A}}^{\text{-}}\right]}$

Henderson-Hasselbalch equation explains the relationship between $pH$ of solution and $p{K}_a$ of acid. For a dissociation of acid $\left(HA\right)$ in aqueous solution,

$HA+H_{2}O\underset{}{\overset{}{\rightleftharpoons }}{H}_3{O}^{+}+A^{-}$

$p{K}_a=pH+\log \frac{\left[HA\right]}{\left[A^{-}\right]}$

During a dissociation of acid in aqueous solution,

• If $pH=p{K}_a$, the concentration of compound in its acidic and basic form is equal.
• If $pH<p{K}_a$, the compound exist in its acidic form.
• If $pH>p{K}_a$, the compound exist in its basic form.

Buffer solutions: Buffer solution is defined as a solution that oppose changes in $\text{pH}$ while adding little amount of either an acid or a base. Buffer solution is a combination of a weak acid and its conjugate base or a weak base and its conjugate acid.

$K_a$: Acid dissociation constant ${\text{K}}_{\text{a}}$ represents how strong the acid is in a solution.

$p{K}_a=-\log \left[K_a\right]$ It is negative logarithm of acid dissociation constant ${\text{K}}_{\text{a}}$.