Chapter 10, Problem 10.103CP

(a)

Interpretation Introduction

Interpretation:

The Henderson-Hasselbalch equation for the given buffer system should be identified.

Concept Introduction:

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.

Henderson-Hasselbalch equation

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

$\text{pH}\text{=}{\text{pK}}_{\text{a}}\text{+}\text{log}\frac{\text{[conjugate base]}}{\text{[acid]}}$

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.

$\text{pH}$: 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.

$K_a$: 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 ${\text{K}}_{\text{a}}$.

(b)

Interpretation Introduction

Interpretation:

The ratio of $HP{O}_4^{2-}$ to $H_{2}P{O}_4^{-}$ require to maintain $pH$ of blood 7.40 should be determined.

Concept Introduction:

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.

Henderson-Hasselbalch equation

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

$\text{pH}\text{=}{\text{pK}}_{\text{a}}\text{+}\text{log}\frac{\text{[conjugate base]}}{\text{[acid]}}$

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.

$\text{pH}$: 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.

$K_a$: 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 ${\text{K}}_{\text{a}}$.