Final answer:
Standard free energy changes cannot be used to determine reaction spontaneity under nonstandard conditions, measure absolute free energies or enthalpies, and cannot violate the restrictions imposed by the second law of thermodynamics such as transferring energy without work.
Step-by-step explanation:
With standard free energy and other standardized values, it is important to note what you cannot do. Standard free energy change (ΔG°) is the change in free energy when substances in their standard states are converted to other substances, also in their standard states. You cannot use ΔG° to directly determine the spontaneity of reactions under nonstandard conditions because these values are calculated using standard pressure (1 bar) and concentrations (1 M), which do not necessarily represent the actual conditions during a reaction. Moreover, you cannot measure absolute free energies or enthalpies; only changes in these quantities can be measured. This limitation stems from the impossibility of defining an absolute zero for enthalpy or Gibbs free energy.
Furthermore, the second law of thermodynamics imposes limitations that prevent energy from being transferred arbitrarily. For example, you cannot unmix substances or convert internal energy into work without accounting for the surroundings and without inputting additional energy. The interpretations of standard Gibbs energy changes should predominantly serve as a rough guide to the potential for spontaneity of a reaction under standard conditions.
Lastly, standard free-energy changes should not be confused with "standard temperature and pressure" (STP) used in gas law calculations, as they refer to different sets of conditions.