The discussion clarifies the distinction between ΔG (Gibbs free energy change) and ΔG° (standard Gibbs free energy change) in thermodynamics. ΔG° represents the free energy change when all reactants and products are in their standard states, while ΔG accounts for the actual conditions of the reaction, expressed through the equation ΔG = ΔG° + RTlnQ. At equilibrium, where the reaction quotient Q equals the equilibrium constant Keq, ΔG becomes zero, leading to the relationship ΔG° + RTlnKeq = 0.
PREREQUISITESChemistry students, educators, and professionals in thermodynamics or chemical engineering seeking a deeper understanding of Gibbs free energy concepts and their applications in reaction dynamics.
Yes I have come across this, but didn't realize the exact difference. Thank you for the clarification :Dmjc123 said:ΔG° is the value of ΔG with all reagents in their standard states. You ought to have encountered this definition in your study of thermodynamics.
For example, if ΔG is the free energy change for a reaction in a system described by the reaction quotient Q (have you come across this?), then
ΔG = ΔG° + RTlnQ
At equilibrium Q = Keq and ΔG = 0, so
ΔG° + RTlnKeq = 0