Abstract
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.
-
Author contributions: All the authors have accepted responsibility for the entire content of this submitted manuscript and approved submission.
-
Research funding: None declared.
-
Conflict of interest statement: The authors declare no conflicts of interest regarding this article.
References
[1] S. R. De Groot and P. Mazur, Non-Equilibrium Thermodynamics, New York, Dover, 1985.Search in Google Scholar
[2] R. H. Petrucci, F. G. Herring, J. D. Madura, and C. Bissonnette, General Chemistry: Principles and Modern Applications, 11th ed. Toronto, Pearson College, 2016.Search in Google Scholar
[3] J. C. Kotz, P. M. Treichel, J. R. Townsend, and D. A. Treichel, Chemistry & Chemical Reactivity, 9th ed. Boston, MA, USA, Cengage Learning Brooks/Cole, 2015.Search in Google Scholar
[4] I. N. Levine, Physical Chemistry, 6th ed. New York, McGraw-Hill, 2009.Search in Google Scholar
[5] W. B. Jensen, “Generalizing the phase rule,” J. Chem. Educ., vol. 78, p. 1369, 2001. https://doi.org/10.1021/ed078p1369.Search in Google Scholar
[6] M. Hillert, Phase Equilibria, Phase Diagrams and Phase Transformations: Their Thermodynamic Basis, 1st ed. Cambridge, Cambridge University Press, 1998.Search in Google Scholar
[7] J. P. Tomba, “Understanding chemical equilibrium: the role of gas phases and mixing contributions in the minimum of free energy plots,” J. Chem. Educ., vol. 94, p. 327, 2017. https://doi.org/10.1021/acs.jchemed.6b00726.Search in Google Scholar
[8] D. Kingston and A. C. Razzitte, “Entropy production in chemical reactors,” J. Non-Equilibrium Thermodyn., vol. 42, p. 265, 2017. https://doi.org/10.1515/jnet-2016-0066.Search in Google Scholar
[9] W. Muschik, “Internal variables in non-equilibrium thermodynamics,” J. Non-Equilibrium Thermodyn., vol. 15, p. 127, 1990. https://doi.org/10.1515/jnet.1990.15.2.127.Search in Google Scholar
[10] K. Denbigh, The Principles of Chemical Equilibrium, 4th ed. Cambridge, Cambridge University Press, 1981.10.1017/CBO9781139167604Search in Google Scholar
[11] R. DeHoff, Thermodynamic in Materials Science, 2nd ed. Boca Raton, CRC Press, Taylor & Francis, 2006.Search in Google Scholar
[12] The sign convention recommended by International Union for Pure and Applied Chemistry is used in this paper. Note that the opposite convention (A = ΔrG) is used in a number of books and papers.Search in Google Scholar
© 2022 Walter de Gruyter GmbH, Berlin/Boston
