In chapter 4 we defined solubility as the amount of a substance that dissolves in a given quantity of solvent at a given temperature to form a saturated solution. We approached solubility in a very simple manner by classifying salts as soluble or insoluble and solubility guidelines were established indicating if a precipitate would form when two solution are combined.
17.4 Overview of Solubility
Experimentally, we can consider solubility equilibria to make predictions about the amount of a given salt that will dissolve. First, we need to define what we mean by solubility and we need to consider what we already know about solubility and compare that to what a chemist thinks about when they hear the term solubility. When chemists discuss solubility they do so in terms of the solubility product constant. Keep in mind that the solubility product constant and molar solubility are two completely different terms. As chemists, we are able to calculate one from the other, but directly comparing the two without a calculation can lead to some improper assumptions. It is very important to recognize the similarities/differences between solubility in terms of the amount (in grams) that dissolve per liter of solution, the molar solubility, and the Ksp.
17.4 The Solubility Product Constant
Example Problem: 1.4 x 10-6 grams of ZnCO3 dissolve in 1.000 mL of solution and 2.8 x 10-6 grams of BaCrO4 dissolve in 1.000 mL of solution, which one has the largest Ksp? Solution: They are the same. Click below to see the explanation.
One common mistake students make is to simply look at the Ksp and think that constant tells them everything they need to know about the solubility when two slightly soluble salts are compared. If you are thinking this way, then you certainly aren’t thinking like a chemist. A chemist would think through the process shown in the example below.
17.4 Ranking the solubility of slightly soluble salts given the Ksp Part I
17.4 Ranking the solubility of slightly soluble salts given the Ksp Part II