Chapter 19: Ionic Equilibria in Aqueous Systems

Students learn the common-ion effect and how to calculate buffer capacity using the Henderson-Hasselbalch equation. Practical examples, including acetic acid/acetate and ammonia/ammonium buffers, show how biological and laboratory systems maintain stable pH. The chapter then covers titration curves, beginning with strong acid–strong base systems before moving to weak acid–strong base, weak base–strong acid, and polyprotic acid titrations. Students learn how to interpret curves, identify equivalence points, and apply acid-base indicators. The next section introduces solubility equilibria, explaining the solubility product constant (Ksp) and how to calculate solubility from Ksp values. The effects of common ions and pH on solubility are discussed, along with real-world applications such as the formation of kidney stones and environmental precipitation processes. The chapter also explains how to predict precipitation reactions and use selective precipitation to separate ions in mixtures. Finally, the focus shifts to complex ion equilibria, where ligands bind to central metal ions, dramatically altering solubility and stability. Examples include the formation of silver-ammonia complexes and amphoteric hydroxide ions. By the end of the chapter, students understand how ionic equilibria govern solubility, buffering, titration behavior, and complex ion stability, all of which are central to biological systems, industrial processes, and environmental chemistry.