Solutions

SOLUTIONS, optically and chemically homogeneous liquids consisting of two (or a greater number) of different kinds of molecules. If each of the components is a liquid and both are present in approximately equal amounts, so that there is no basis for calling one of them the solvent and the other the dissolved substance, one usually does not speak of solutions, but of liquid mixtures (similarly to how one speaks of a mixture of gases, not of a solution of one gas in another). Thus, a sharp boundary cannot be established between solutions and mixtures. If the mixture is a solid body, it is significantly more difficult to obtain an even distribution of one component in another. If this is nevertheless achieved, one speaks of a 'solid solution'. The degree of dispersion of the dissolved substance is also an essential circumstance. Only systems in which the dissolved substance is distributed in the form of individual molecules (simple or associated) or the ions formed by them are considered true solutions. In cases where it forms larger aggregates of ultramicroscopic dimensions, one speaks of colloidal solutions or sols (see). However, simple mechanical dispersion of the dissolved substance to molecular dimensions is not sufficient for the formation of solutions. Apparently, a necessary condition for this is also a certain solvation (in the case of aqueous solutions - hydration) of the dissolved molecules, a certain affinity between them and the solvent. The basis for the modern theory of solutions was provided by the research of van't Hoff. Van't Hoff drew attention to the remarkable analogy between the pressure of gases and the osmotic pressure produced by a solution contained in a semipermeable osmotic cell (see Osmotic pressure). The laws established by Pfeffer and de Vries for solutions exactly correspond to the gas laws. This led van't Hoff to formulate the basic principle of the theory of solutions: 'the osmotic pressure of a solution equals the pressure that the dissolved substance, in the molecular state, would produce as a gas or vapor in the same volume and at the same temperature'. Proceeding from this, it was natural to extend to solutions the kinetic theory developed for gases (see). The latter explains the laws of ideal gases on the basis of the concept of the chaotic motion of point molecules, not connected by any forces except the forces of elastic impact when they collide with each other. Similarly, the kinetic theory explains the properties of solutions and the osmotic pressure they create. However, for both gases and solutions, the kinetic theory in its simplest form gives satisfactory results only under certain conditions, in particular at sufficiently low concentration of gas or dissolved substance. Ideal dilute gases correspond to ideal dilute solutions. At higher concentrations, the inaccuracy of both basic premises of the theory becomes more pronounced: the concept of molecules as material points (and therefore not having definite, finite dimensions) and the absence of forces of interaction between them. Van der Waals, abandoning both these simplifications, taking into account the volume occupied by the molecules themselves and the forces of attraction acting between them (which decrease extremely rapidly with distance), built the theory of real gases and derived an equation quantitatively expressing their behavior. This equation is also applicable to solutions, and as their concentration increases, the correction for the volume occupied by the dissolved molecules themselves becomes increasingly important: the volume remaining free for molecular motion decreases by a corresponding amount. In practice, this correction is often introduced in such a way that the concentration of the dissolved substance is related not to the entire volume of the resulting solution, but to the mass of the solvent alone, which can be approximately equated to the space remaining free for the movement of dissolved molecules. Unlike the usual method of calculating concentrations based on the volume of the solution, this method of calculation based on the weight of the solvent is mostly applied to highly concentrated solutions. Further complications arise when applying the theory to electrolyte solutions. The deviations observed in them were explained by Arrhenius within the framework of the old kinetic theory. The dissociation of molecules into ions accepted by him leads only to an increase in the concentration of freely moving particles without any influence on the forces acting between them (see Electrolytic dissociation). Only the gradual clarification of the insufficiency of Arrhenius' theory for explaining the behavior of strong electrolytes led in the last twenty years to the construction of a new theory of ionic solutions - the theory of activity. Its most essential feature is the recognition and consistent accounting of those electrostatic interionic forces that inevitably arise in electrolyte solutions as a result of the most characteristic property of the ion - the presence of a free electric charge in it. These forces acquire serious significance at significantly lower concentrations than Van der Waals forces of attraction. They affect the most fundamental property of the ion from the point of view of kinetic theory - its molecular kinetic energy (and consequently all properties that depend on its motion). Arrhenius, Kohlrausch and all researchers who stood on the basis of the old theory assumed that all independently moving particles (whether ions or molecules) have on average the same kinetic energy, proportional to the absolute temperature. All changes in the magnitude of osmotic pressure (and the associated changes in vapor pressure, freezing point and boiling point), on the one hand, and electrical conductivity, on the other, were therefore reduced exclusively to changes in the number of freely mobile particles and could therefore be used to measure the degree of dissociation. In contrast to this, Bjerrum, Milner, Debye and Huckel showed that the kinetic energy and mobility of ions themselves are functions of the concentration of ionic charges contained in the solution. It decreases at constant temperature as the total ionic concentration increases. Therefore, changes in osmotic pressure, electrical conductivity and chemical activity of the electrolyte cannot serve as an indicator of proportional changes in the concentration of ions in the solution. Furthermore, Arrhenius' theory allowed one to characterize all these changes with one quantity - the degree of dissociation: the concentration of free ions determined both the electrical conductivity of the solution and the active concentration of the ion in chemical reactions and the increase in the theoretical value of osmotic pressure it produces. New concepts force the introduction of special activity coefficients to characterize the participation of the ion in each of these processes. The same change in the average speed of the ion has different effects on them: for electrical conductivity, the mobility of the ion is decisive, for osmotic pressure - its kinetic energy, while chemical activity depends both on the number of molecular collisions and their energy. The concept of interionic electrostatic forces as a factor playing - at least in the case of solutions of strong electrolytes - a decisive role, shifts the center of gravity of the doctrine of ionic equilibria from chemistry to the realm of physics. Indeed, if the concentration of an ion can be changed by chemical binding, its activity is entirely determined by physical conditions, the most important of which, besides temperature, is the electrostatic field created by all ions present in the solutions, regardless of their chemical nature.