Chapter 2: Water, pH, & Acid–Base Balance

Water, pH, & Acid–Base Balance explains how the dipolar structure of water and its ability to form hydrogen bonds facilitate its function as an ideal solvent, possessing a high dielectric constant that reduces the attraction between charged species like salts. The text details the various noncovalent forces that stabilize macromolecular structures, including salt bridges, van der Waals forces, and hydrophobic interactions, the latter of which are governed by entropic changes that dictate the folding of proteins and the formation of lipid bilayers. Water is also characterized as a potent nucleophile involved in metabolic hydrolysis, with enzymes playing a pivotal role in accelerating these reactions or sequestering substrates to allow for biopolymer synthesis in an aqueous environment. The discussion transitions into the quantitative aspects of acidity, defining pH as the negative logarithm of the hydrogen ion concentration and explaining how the Henderson-Hasselbalch equation, expressed as pH = pKa + log [Conjugate Base] / [Acid], describes the behavior of weak acids. Clinical relevance is highlighted through the exploration of acid-base balance, where acidosis occurs at a blood pH lesser than 7.35 and alkalosis at a pH greater than 7.45. Finally, the chapter outlines the critical function of physiological buffering systems, such as bicarbonate, orthophosphate, and proteins, which maintain homeostatic pH levels by resisting significant changes when protons are added or removed from the system.