Solubility
The extent to which a solute will dissolve in a solvent is called its solubility. The solubility of a chemical is conventionally expressed as the maximum number of grams of a chemical that will dissolve in l00g of solvent but conversion to mol L-1 or g L-1 is simple and may be appropriate for some applications (see below). Since solubility is temperature dependent, is always quoted at a specific temperature. With a very few exceptions, increasing the temperature of a solvent increases the solubility of the solute.Saturated solutions
For practical purposes, a saturated solution is one in which no more solute will dissolve. For example, the solubility of sodium chloride in water is 35.6g per l00g at 25°Cand 39.1 g per 100g at 100°Cand both solutions are saturated solutions at their respective temperatures. If the 100°C solution is cooled to 25°C, then 3.5 g of NaCl crystals will precipitate from the solution, because the solution at 25°C requires only 35.6 g of NaCI for saturation. This process is the basis of purification of compounds by recrystallization.
Solubility product In dilute aqueous solutions, it has been demonstrated experimentally for poorly soluble ionic salts (solubilities less than 0.01 mol L-1) that the mathematical product of the total molar concentrations of the component ions is a constant at constant temperature. This product, Ks is called the solubility product. Thus for a saturated solution of a simple ionic compound AB in water, we have the dynamic equilibrium:
ABsolid ↔ A+(aq) + B−(aq)
where AB represents the solid which has not dissolved, in equilibrium with its ions in the aqueous saturated solution. Then:
Ks = [A+] × [B−]
For example, silver chloride is a solid of solubility 0.000 15g per 100mL of water in equilibrium with silver cations and chloride ions. Then:
Ks = [Ag+] × [Cl−]
The solubility of AgCl is 0.0015gmL−1 (l0 × solubility per 100g, assuming that the density of water is 1.0gmL−1) and therefore the solubility of AgCl is 0.0015 ÷ 143.5 = 1.05 × 10−5 mol L−1. Thus the saturated solution contains 1.05 × 10−5 mol L−1 of Ag+ ions and 1.05 x 10−5 mol L−1 of Cl− ions and the solubility product Ks is:
Ks = (1.05 × 10−5) × (1.05 × 10−5) = 1.1 × 10−10 mol2 L−2
If the solid does not have a simple I: I ratio of its ionic components, e.g. PbCb, then the solubility product is given by:
Ks = [Pb2+] × [Cl−]2
In general terms, the solubility product for a compound MyNx, is given by:
Ks = [M+]y × [N−]z
The practical effects of solubility products are demonstrated in the detection of anions and cations by precipitation and in quantitative gravimetric analysis (p. 139). For example, if dilute aqueous solutions of silver nitrate (solubility 55.6 g per 100g of water) and sodium chloride (solubility 35.6 g per 100g of water) are mixed, an immediate white precipitate of AgCl is produced because the solubility product of AgCI has been exceeded by the numbers of Ag+ and Cl− ions in the solution, even though the ions come from different 'molecules'. A saturated solution of AgCl is formed and the excess AgCI precipitates out. The solubility product of the other combinations of ions is not exceeded and thus sodium and nitrate ions remain in solution. Even if the concentration of Ag+ is extremely low, the solubility product for AgCl can be exceeded by the addition of an excess of Cl− ions, since it is the multiplication of these two concentrations which defines the solubility product. Thus soluble chlorides can be used to detect the presence of Ag+ ions and, conversely, soluble silver salts can be used to detect Cl− ions, both quantitatively and qualitatively.
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