In the recent article by Craig, I was pleased to find an excellent critique of the use of deterministic terms in thermodynamics and physical chemistry (1). Chemical reactions have a probabilistic nature; so, according to the second law of thermodynamics, their direction is always determined by the sign of the total entropy change:
∆Stotal = ∆Ssystem + ∆Ssurroundings |
(1) |
I would like to indicate one particularly misleading misconception that arises early in general chemistry, while teaching simple exchange reactions in aqueous solutions. A simple and easy-to-remember ‘explanation’, that precipitation of a solid and/or formation of water are “driving forces” of those reactions or “drive them to completion”, still occurs among instructors. This is a disservice to students, particularly when they learn that the formation of solid precipitates as well as combination of two ions, H+(aq) and OH–(aq), to form one particle, H2O(l), are supposed to drive reactions backwards because they are unfavorable in terms of entropy.
Given the small negative ∆H° [–55.7 kJ/mol (2)] and much more negative ∆G° (based on Kw = 10–14), the ∆S° of neutralization at room temperature is actually positive (unlike that in the gas phase) because of “unfreezing” of water molecules participating in ions’ hydration in aqueous solutions. For the same reason [significant hydration of the reacting ions (3)], ∆Ssystem is also positive for precipitation reactions in aqueous solutions at room temperature (4).
In addition, neutralization and some (although not all) precipitation reactions are exothermic (2, 4). As a result, a large ∆Ssurroundings = ∆Hsurroundings/T term contributes to the positive sign of ∆Stotal in eq 1. It is the combination of a negative ∆Hsystem and positive ∆Ssystem at room temperature that creates an illusion of these reactions’ “completion” (of course, they are never complete). However, at elevated temperatures poor hydration of the ions (reactants) forces the switch of the sign of ∆Ssystem. Thus, at high temperatures, at low reactant concentrations, and/or for weaker acids/bases the opposite reactions, i.e., hydrolysis of salts and dissolution of precipitates, may prevail.
Literature Cited
- Craig, N. C. J. Chem. Educ. 2005, 82, 827–828.
- Bertrand, G. L.; Millero, F. J.; Wu, C.; Hepler, L. G. J. Phys. Chem. 1966, 70, 699–705.
- Marcus, Y. Biophysical Chemistry 1994, 51, 111–127.
- Wagman, D. D.; Kilday, M. V. J. Res. National Bureau of Standards, A: Physics and Chemistry 1973, 77, 569–579; Evans, W. J.; Marini, M. A.; Martin, C. J. J. Inorg. Biochem. 1983, 19, 129–132.
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