Chapter 5: Simple Mixtures

Simple Mixtures extends fundamental thermodynamic principles to analyze the behavior of simple mixtures and solutions. At the core of mixture thermodynamics lies the concept of chemical potential, which represents the Gibbs energy contribution of a single component and must achieve equilibrium across all phases present in a system. For ideal solutions, the chemical potential varies logarithmically with concentration, a relationship quantified through Raoult's law for volatile components and Henry's law for dilute solutions where one component dominates. The thermodynamic consequences of mixing are examined through enthalpy, entropy, and Gibbs energy changes, revealing that ideal solutions exhibit no enthalpy of mixing yet generate negative, spontaneous Gibbs energy changes driven primarily by entropy increase. The chapter then addresses colligative properties, which depend solely on solute particle count rather than chemical identity, including vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure. These properties provide experimental methods for determining the molar masses of large molecules such as proteins. The behavior of complex mixtures is systematized through phase diagrams, beginning with binary systems that display liquid-liquid equilibrium phenomena like azeotropic points and solid-liquid behavior including eutectic points and incongruent melting reactions. Ternary phase diagrams extend this analysis to three-component systems using triangular representations. Because real solutions deviate significantly from ideal behavior due to intermolecular forces and molecular size effects, the chapter introduces activity as an effective concentration that accounts for these deviations. For ionic solutions where strong electrostatic interactions dominate, the Debye-Hückel limiting law provides a theoretical framework for calculating mean activity coefficients and predicting thermodynamic properties of electrolytes at low concentrations.