The equation ΔG^∘=-R T lnK relates the value of Kp, not Kc, to ΔG^* for gas-phase reactions. Exp

step 1 When we define the K constant, we are considering the activity of each of the species in the rea

The equation ΔG^∘=-R T lnK relates the value of Kp, not Kc,...

Key Concepts

Gibbs Free Energy and Equilibrium

The equation ?G° = –RT ln K links the standard Gibbs free energy change to the equilibrium constant. This relation is derived using the activities of the reactants and products under standard state conditions. For gas-phase reactions, since activities are defined in terms of pressure, K in this equation naturally corresponds to Kp rather than Kc.

Standard States in Thermodynamics

Standard states establish reference conditions for thermodynamic measurements. In the case of gases, the standard state is typically defined by a pressure of 1 bar (or 1 atm in some conventions), rather than a concentration. This definition is central when relating Gibbs free energy changes to equilibrium constants because it determines the form of the activities used in these calculations.

Thermodynamic Activity

Activity is a dimensionless quantity representing the 'effective concentration' of a species in a mixture. For gases, the activity is defined as the ratio of its partial pressure to a standard pressure. This is why the equilibrium constant in the ?G° = –RT ln K equation is expressed in terms of partial pressures (Kp) for gas-phase reactions rather than molar concentrations (Kc).

Equilibrium Constants (Kp vs Kc)

Equilibrium constants can be defined based on either pressures (Kp) or concentrations (Kc). In gas-phase reactions, Kp is more appropriate because it directly uses partial pressures in accordance with the standard state of gases. Converting between Kp and Kc typically requires using the ideal gas law, which introduces temperature and volume factors.

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