Chapter 17: Spontaneous Change: How Far?

The central distinction separates thermodynamic favorability from reaction kinetics: a process may be thermodynamically driven to occur yet proceed at imperceptibly slow rates due to kinetic barriers. The chapter grounds spontaneity in entropy and the second law of thermodynamics, which states that any spontaneous process increases the total entropy of the universe. Entropy itself represents the dispersal of energy and matter within a system, with the Boltzmann perspective linking entropy to the number of accessible microstates. Standard molar entropy values generally increase from solid to liquid to gas states and rise with molecular complexity and temperature. To avoid calculating universe-scale entropy changes, chemists employ Gibbs free energy, defined through the relationship G equals H minus TS, which integrates both enthalpy and entropy contributions into a single criterion. At constant temperature and pressure, spontaneous change occurs when free energy decreases, with equilibrium reached when free energy change equals zero. The chapter develops mathematical connections linking standard free energy change to the equilibrium constant through the fundamental equation ΔrG° equals negative RT ln K, and to electrochemical cell potentials through ΔrG° equals negative nFE°cell. The reaction quotient enables prediction of spontaneity under non-standard conditions using ΔrG equals ΔrG° plus RT ln Q. Temperature significantly influences both equilibrium constants and spontaneity through the van't Hoff equation, which relates equilibrium constants at different temperatures using the standard enthalpy change. The chapter concludes with a practical application examining photochemical smog formation, demonstrating how high-temperature combustion shifts equilibrium toward nitrogen oxide production, which subsequently participates in atmospheric reactions driven by solar energy to generate secondary pollutants such as ground-level ozone and peroxyacetylnitrate.