![]() ![]() ![]() If the concentrations are such that Q is less than K sp, then the solution is not saturated and no precipitate will form. A saturated solution in equilibrium with the undissolved solid will result. The reaction shifts to the left and the concentrations of the ions are reduced by formation of the solid until the value of Q equals K sp. If we mix a solution of calcium nitrate, which contains Ca 2+ ions, with a solution of sodium carbonate, which contains CO 3 2- ions, the slightly soluble ionic solid CaCO 3 will precipitate, provided that the concentrations of Ca 2+ and CO 3 2- ions are such that Q is greater than K sp for the mixture. If we add calcium carbonate to water, the solid will dissolve until the concentrations are such that the value of the reaction quotient (Q = ) is equal to the solubility product ( K sp = 8.7 × 10 –9). We can establish this equilibrium either by adding solid calcium carbonate to water or by mixing a solution that contains calcium ions with a solution that contains carbonate ions. However, when we add an excess of solid AgCl to water, it dissolves to a small extent and produces a mixture consisting of a very dilute solution of Ag + and Cl – ions in equilibrium with undissolved silver chloride: Recall from the solubility rules in an earlier chapter that halides of Ag + are not normally soluble. Silver chloride is what’s known as a sparingly soluble ionic solid (Figure 17.6a). We will also learn how to use the equilibrium constant of the reaction to determine the concentration of ions present in a solution. In this section, we will find out how we can control the dissolution of a slightly soluble ionic solid by the application of Le Châtelier’s principle. ![]() We want the calcium carbonate in a chewable antacid to dissolve because the CO 3 2- ions produced in this process help soothe an upset stomach. On the other hand, sometimes we want a substance to dissolve. Preventing the dissolution prevents the decay. The dissolution process is aided when bacteria in our mouths feast on the sugars in our diets to produce lactic acid, which reacts with the hydroxide ions in the calcium hydroxylapatite. Tooth decay, for example, occurs when the calcium hydroxylapatite, which has the formula Ca 5(PO 4) 3(OH), in our teeth dissolves. In some cases, we want to prevent dissolution from occurring. The preservation of medical laboratory blood samples, mining of sea water for magnesium, formulation of over-the-counter medicines such as Milk of Magnesia and antacids, and treating the presence of hard water in your home’s water supply are just a few of the many tasks that involve controlling the equilibrium between a slightly soluble ionic solid and an aqueous solution of its ions. Carry out equilibrium computations involving solubility, equilibrium expressions, and solute concentrations.Write chemical equations and equilibrium expressions representing solubility equilibria. ![]()
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