Chapter 12: Reactions in Pure Solids & Gas Phases

Reactions in Pure Solids & Gas Phases establishes the fundamental thermodynamic principles necessary for analyzing chemical reactions that occur between pure condensed phases (solids or liquids) and a gaseous phase, such as the oxidation of a pure metal to its pure oxide. A major thermodynamic simplification is defining the standard state for a pure solid or liquid simply as the pure species at temperature T, without specifying pressure, because the Gibbs free energy of condensed phases is highly insensitive to pressure variations in the usual range of pressures (0 to 1 atm). For a system to achieve complete equilibrium, phase equilibrium must be established concurrently with reaction equilibrium. Consequently, the equilibrium constant (K) for the overall reaction can often be written exclusively in terms of the partial pressures of the species that exist only in the gas phase. The system’s spontaneity (whether oxidation or reduction occurs) is determined by comparing the actual partial pressure of the gaseous reactant (e.g., oxygen) to the unique equilibrium partial pressure (p oxygen equilibrium) that is fixed solely by the system temperature. Applying the Gibbs Equilibrium Phase Rule, systems composed of two pure solids and one gas phase are found to maintain only one degree of freedom, meaning that the equilibrium gas pressure is uniquely fixed once the temperature is set. The Ellingham diagram is introduced as a critical analytical tool, plotting the standard Gibbs free energy change (Delta G standard) versus absolute temperature. These lines are approximated as straight over ranges of temperature where no phase change occurs, with the slope corresponding to the negative standard entropy change (minus Delta S standard) and the intercept relating to the standard enthalpy change. Since many metal oxidation reactions involve the disappearance of one mole of oxygen gas, their standard entropy changes are similar, causing the Ellingham lines for different metals to run roughly parallel. A more negative value of Delta G standard indicates a greater chemical affinity between the metal and oxygen, signifying a more stable oxide. Phase transformations, such as the melting of a metal or an oxide, introduce an abrupt upward or downward change in the line's slope, referred to as an "elbow," because they alter the standard enthalpy and standard entropy of the reaction. The utility of the diagram is enhanced by the Richardson nomographic scales, which allow for the geometric determination of the unique equilibrium partial pressures (p oxygen equilibrium) or gas ratios (such as p carbon monoxide divided by p carbon dioxide or p hydrogen divided by p water) necessary for equilibrium at any given temperature. Intersections between stability lines compare the relative stability of different oxides, showing the minimum temperature (T sub E) required for one metal to reduce another’s oxide. The chapter also details the oxides of carbon, examining the vital Boudouard reaction (carbon plus carbon dioxide gives two carbon monoxide), which is essential for reduction metallurgy, often visualized through plots comparing the logarithmic ratio of p carbon dioxide divided by p carbon monoxide versus temperature to delineate stability regions.