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Among the chemical reactions having a pronounced thermodynamic driving force to the formation of products, a distinction has to be made between processes which attain a chemical equilibrium state very much shifted towards the products and those where the final equilibrium state corresponds to the truly complete consumption of reactant(s), i.e. where a true chemical equilibrium is actually not attained under the given conditions. Based on a few selected examples, two thermodynamic arguments are led which rationalise the above distinction from different points of view: a phase-rule point of view and a Gibbs energy minimization approach. “Conceptually complete” reactions involve pure phases and, as a consequence, establishing chemical equilibrium would imply a negative variance, what is avoided by the complete consumption of one or more phases. In the Gibbs energy approach, “conceptually complete” reactions and “practically complete” ones can be distinguished (at fixed temperature and pressure) by the different way to attain the minimum Gibbs energy condition, respectively with sharp (not differentiable) and flat (zero derivative) minimum points as a function of the extent of reaction  ξ .

Are all chemical reactions in principle reversible? Thermodynamic distinction between “conceptually complete” and “practically complete” reactions