Thermodynamics fundamentals: Pick the correct statement among the following options.

Introduction / Context:
Accurately distinguishing thermodynamic quantities such as ΔG, ΔH, and ΔE and recognizing system classifications is crucial. This question checks conceptual clarity on equilibrium criteria, sign conventions, and system boundaries—concepts that underpin reaction feasibility and process design.

Given Data / Assumptions:

• Standard definitions of open, closed, and isolated systems.
• Equilibrium at constant temperature and pressure (common in chemistry and process engineering).
• Heats of solution can be exothermic or endothermic depending on solute–solvent interactions.

Concept / Approach:
The fundamental criterion for equilibrium at constant T and P is ΔG = 0. Spontaneous change at constant T and P has ΔG < 0, while non-spontaneous changes have ΔG > 0. Heats of solution vary in sign based on dissolution energetics. The classification of systems depends on exchange of mass and energy: an open beaker exchanging heat with surroundings and evaporating vapor is not isolated.

Step-by-Step Solution:

Verification / Alternative check:
From the Gibbs criterion, chemical potentials of all species are equalized at equilibrium and ΔG for any infinitesimal change is zero.

Why Other Options Are Wrong:

• (a) Overgeneralizes dissolution energetics.
• (c) Misapplies thermodynamic relations.
• (d) Misclassifies the system; isolated systems exchange neither mass nor energy.

Common Pitfalls:
Confusing ΔG with ΔE or ΔH; overlooking that “isolated” forbids both heat and mass transfer.

Final Answer:
At thermodynamic equilibrium, the Gibbs free energy change ΔG is zero.