Скачать или смотреть PH and Buffers lecture 2- APPLICATION OF BUFFERS, BUFFER EQUATION

This video focuses on the Applications and Theory of Buffers, specifically within pharmaceutical and biological systems.
Core Definitions
Buffers: Compounds or mixtures that resist changes in pH when small amounts of acid or alkali are added.
Buffer Action: The actual resistance to pH change shown by a buffer.
Buffer Capacity: A quantitative measure of how much acid or base is required to change the pH of a solution by one unit.
Pharmaceutical Applications
The document highlights why pH control is essential in drug formulation:
Solubility: Maintaining a specific pH ensures the drug remains dissolved and does not precipitate during its shelf-life.
Stability: Adjusting pH prevents chemical degradation, such as hydrolysis.
Biological Activity: Enzymes and certain active ingredients only function optimally at specific pH values.
Patient Comfort: Formulations for eyes (lacrimal fluid), blood (injections), or abraded skin must match physiological pH to avoid irritation, burning, or tissue damage (hemolysis).
Biological Buffer Systems
The document details how the human body maintains pH balance:
Blood: Maintains a strict pH of approximately 7.4.
Primary Buffers (Plasma): Carbonic acid-bicarbonate system and phosphoric acid salts.
Secondary Buffers (Erythrocytes): Hemoglobin/oxyhemoglobin and phosphoric acid salts.
Lacrimal Fluid (Tears): Has a high buffer capacity with a pH of about 7.4.
Urine: Typically has a pH of 6.0, but can range from 4.5 to 7.8 depending on the body's need to excrete hydrogen ions.
Types of Buffer Systems
Weak acid + conjugate base (e.g., Acetic acid and Sodium acetate).
Weak base + conjugate acid (e.g., Ammonium hydroxide and Ammonium chloride).
Two salts acting as an acid-base pair (e.g., Potassium phosphate systems).
Amphoteric electrolytes (e.g., Glycine solutions).
Strong acid/base solutions (e.g., Strong HCl with KCl).