Precipitation and Dissolution of Alkaline Earth Sulfates: Kinetics and Surface Energy

Abstract

The precipitation and dissolution of sparingly soluble alkaline earth sulfates play important roles in the establishment of geological equilibria (Davis and Hayes 1986). There is little doubt that there are close relationships among the interfacial tensions between the solid phases and their solutions, the observed solubilities, and the kinetics of crystallization and dissolution. The experimental methods used for investigating the rates of these reactions are reviewed with particular reference to the Constant Composition (CC) technique (Tomson and Nancollas 1978) which enables both crystallization and dissolution reactions to be studied at constant thermodynamic driving forces. Interfacial tensions of the alkaline earth sulfates may be measured using thin-layer wicking methods (Van Oss et al. 1992), and surface-tension components theory yields detailed information concerning the Lifshitz–van der Waals and Lewis acid–base parameters of the surface tension. The mechanisms of the growth and dissolution reactions are discussed from the point of view of kinetics and interfacial energies. ### Supersaturation Studies of crystallization from solution first require the preparation of metastable, supersaturated solutions in which the concentration of solute exceeds the saturation value at a given temperature and pressure. Some common expressions describing supersaturation are defined by Equations (1)–(3): \batchmode \documentclass[fleqn,10pt,legalpaper]{article} \usepackage{amssymb} \usepackage{amsfonts} \usepackage{amsmath} \pagestyle{empty} \begin{document} \[{\Delta}C = {-} Cs\] \end{document}(1) \batchmode \documentclass[fleqn,10pt,legalpaper]{article} \usepackage{amssymb} \usepackage{amsfonts} \usepackage{amsmath} \pagestyle{empty} \begin{document} \[\mathit{S} = \frac{C}{C_{S}}\] \end{document}(2) \batchmode \documentclass[fleqn,10pt,legalpaper]{article} \usepackage{amssymb} \usepackage{amsfonts} \usepackage{amsmath} \pagestyle{empty} \begin{document} \[\mathit{{\sigma}} = \frac{C {-} \_{S}}{C\_{S}} = \mathit{S {-}} 1\] \end{document}(3) where C is the solute concentration (moles of solute per liter of solvent, for example), Cs is the solubility value, ΔC is the concentration driving force, S is the supersaturation ratio, and μ is the relative supersaturation. It is important to note that these definitions of supersaturation assume that the solutions are ideal despite the fact that they consist of mixtures of electrolytes and, in many cases, additional nonelectrolytes. The definitions are correct, therefore, only for very dilute solutions for which it is possible to assume that the activity coefficients are unity. In crystallization studies, one inevitably deals with solutions with high concentrations even …