Acid Base Clinical Biochemistry | pH, Ka, pKa, buffers and more!

Introduction to pH and Ion Concentration00:00:00

Dr. Mike introduces the clinical importance of monitoring ion concentrations in blood plasma. He explains why hydrogen ion concentration is represented as pH to avoid calculation errors and highlights the inverse relationship between pH and hydrogen ion concentration.

Acids, Bases, and Dissociation00:10:15

Defines acids as proton donors and bases as proton binders. Introduces the dissociation constant (Ka) as a measure of an acid's strength and its tendency to split apart in a solution.

The Role of Buffers00:18:19

Explains how weak acids function as buffers to resist drastic changes in pH. Establishes that a system acts as an optimal buffer when the pH is equal to the pKa.

The Henderson-Hasselbalch Equation00:20:46

Introduces the Henderson-Hasselbalch equation as the key tool for calculating pH in clinical settings, demonstrating how it relates pH, pKa, and the ratio of conjugate base to acid.

The Bicarbonate Buffer System00:29:14

Explains the bicarbonate buffer system, the body's most important buffer. Dr. Mike demonstrates how to calculate blood pH using partial pressure of CO2 and bicarbonate levels, and how shifts in these values lead to respiratory or metabolic acidosis and alkalosis.