Chapter 7, Problem 10RQ

Summary Introduction

Introduction:

In biochemical reactions, the Henderson–Hasselbalch equation can be used to determine the pH (potential of hydrogen) of a buffer solution, when the concentration of the acid and that of its conjugate base are given.

$\begin{array}{l}\text{H-A}\underset{}{\rightleftharpoons }{\text{H}}^{+}+{\text{A}}^{-}\\ acidbase\\ \end{array}$

$\begin{array}{c}\log Ka=\log [H^{+}]+\log ([A^{-}]/[HA])\\ -pKa=-pH=\log ([A^{-}]/[HA])\\ \text{therefore,}\\ pH=pKa+\log [(ionized/unionized)]\end{array}$

Thus, the Henderson–Hasselbalch equation applied during normal conditions in the blood is,

$\text{pH}=\text{pK}+\log \frac{[{\text{HCO}}_{\text{3}}^{-}]}{[{\text{P}}_{\text{CO2}}\times \text{0}\text{.03]}}$

This equation helps to calculate bicarbonate ion concentration.