Kaplan MCAT General Chemistry Review

Our last topic for this chapter on solutions picks up on a theme that we began discussing in Chapter 5, Chemical Kinetics and Equilibrium. The process of solvation, like other reversible chemical and physical processes, tends toward an equilibrium position defined as the lowest energy state of a system under given conditions of temperature and pressure. Related to determinations of equilibrium is the characterization of processes as spontaneous or nonspontaneous: Systems tend to move spontaneously toward the equilibrium position, but any movement away from equilibrium is nonspontaneous. In the process of creating a solution, the equilibrium is defined as the saturation point and the solute concentration is at its maximum value for the given temperature and pressure. Immediately after solute has been introduced into a solvent, most of the change taking place is dissociation, because no dissolved solute is initially present. However, once solute is dissolved, the reverse process, precipitation of the solute, will begin to occur. When the solution is dilute (unsaturated), the thermodynamically favored process is dissolution, and initially, the rate of dissolution will be greater than the rate of precipitation. As the solution becomes more concentrated and approaches saturation, the rate of dissolution lessens while the rate of precipitation increases. Eventually, the saturation point of the solution is reached, and the solution exists in a state of dynamic equilibrium for which the rates of dissolution and precipitation are equal and the concentration of dissolved solute reaches a steady state (that is, constant) value. Neither dissolution nor precipitation is more thermodynamically favored at equilibrium than the other (because favoring either one of them would necessarily result in the solution no longer being in a state of equilibrium), so the change in free energy is zero, as is the case for all systems at equilibrium.

An ionic solid introduced into a polar solvent dissociates into its component ions, and the dissociation of such a solute in solution may be represented by

A_mB_n (s) mAⁿ+ (aq) + nB^m- (aq)

On Test Day, when you are working through any problem of solution chemistry, the first step you must take is to write out the balanced dissolution equation for the ionic compound in question. This first step is essential for correctly calculating solubility product constant, ion product, molar solubility, or common ion effect. In other words, it is the essential first step for nearly every solution chemistry problem you will see on the MCAT.

THE SOLUBILITY PRODUCT CONSTANT

Most solubility problems on the MCAT deal with solutions of sparingly soluble salts, which are ionic compounds that have very low solubility in aqueous solutions. You may wonder why any ionic compound would not be highly soluble in water. The determination for the degree of solubility is the relative changes in enthalpy and entropy associated with the dissolution of the ionic solute at a given temperature and pressure. One common sparingly soluble salt is silver chloride, AgCl, which dissolves according to the following equation:

AgCl (s) Ag⁺ (aq) + Cl^- (aq)

The law of mass action can be applied to a solution at equilibrium; that is to say, when the solution is saturated and the solute concentration is maximum and dynamically stable. For a saturated solution of the ionic compound with the formula A_mB_n, the equilibrium constant for its solubility in aqueous solution, called the solubility product constant K_sp, can be expressed by

K_sp = [Aⁿ+]^m[B^m-]ⁿ

where the concentrations of the ionic constituents are equilibrium (saturation) concentrations. For example, we can express the K_sp of silver chloride as follows:

K_sp = [Ag⁺] [Cl^-]

You’ll notice that for the law of mass action of solutions, the denominator seems to be missing. Well, if you think back to our discussion of the properties of the equilibrium constant in Chapter 5, you’ll remember that we don’t include the concentration of the pure solids or pure liquids. Since the silver chloride solution was formed by adding pure solid silver chloride to pure water, neither the solid silver chloride nor the water is included.