Advanced Placement (AP) Chemistry Practice Exam

A reaction is considered to be at equilibrium when the Gibbs free energy change (ΔG) equals zero. At this point, the rate of the forward reaction equals the rate of the reverse reaction, meaning that the concentrations of reactants and products remain constant over time. This condition indicates that no net change is occurring in the system, even though molecular activity continues.

When ΔG is negative, it indicates that a reaction is spontaneous and will proceed in the forward direction to reach equilibrium, but the system is not yet at equilibrium. Conversely, when ΔG is positive, the reaction is non-spontaneous and will favor the reverse direction. Both of these scenarios point to a system that has not reached equilibrium. The state of equilibrium is fundamentally characterized by ΔG being zero, signifying that there is no driving force for either direction of the reaction.

ΔH, which represents the change in enthalpy, does not directly indicate equilibrium. A reaction can have a positive, negative, or zero ΔH and still achieve equilibrium based purely on the relationship of reactants and products at a given temperature and pressure. Thus, the critical point for identifying equilibrium is when ΔG equals zero.