95Solubility of Substances in Water
Water is an excellent solvent for many compounds. Some dissolve in it as molecules while others,
called electrolytes, dissociate and dissolve not as neutral molecules but as charged species called
ions. Compounds which exist as solid ionic crystals dissolve in water as ions, and most of them are
highly soluble in water. “Highly soluble” is a somewhat elastic description, but generally means
soluble to at least the extent of forming 0.1 to 1.0 molar aqueous solutions. Salts which are less
soluble in water than this at room temperature are called slightly soluble salts.
The solubility of an ionic salt depends both upon its cations and its anions, but for simple salts in
aqueous solution at room temperature the following general observations are useful. Almost all
sodium, potassium, and ammonium salts are highly soluble; the only significant exception is KCIO 4 ,
which is moderately soluble almost without exception.
Metal carbonates and phosphates are generally insoluble or slightly soluble, with the exception of
those of sodium, potassium, and ammonium which are highly soluble; magnesium ammonium
phosphate is used for the precipitation of magnesium ion.
Metal halides are generally highly soluble, with the exception of those of silver, lead, and mercury
(I). Lead chloride is slightly soluble while silver and mercury (I) chlorides are much less soluble.
Sulfate salts are generally highly soluble as well, with more exceptions; calcium, barium, strontium,
lead, and mercury (I) sulfates are almost insoluble while silver sulfate is slightly soluble. Metal
sulfides are generally insoluble in water.
Solid-Solution (Solubility) Reactions
When solids dissolve, the solutes are no longer pure substances and their activity can no longer
be taken as unity. In dilute solutions, aqueous or otherwise, activities of solutes are often taken as
equal to their molar concentrations. These equilibria are called solubility equilibria and are taken
up under the following main heading. The example below shows how the form in which they are
written compares to other equilibrium constants.
Example. The equilibrium constant for the reaction AgCl(s) <--> Ag+(aq)+Cl-(aq) is written as K=
a(Ag+)a(Cl-)/a(AgCl); more commonly, it is written in the form Ka(AgCl)=a(Ag+)a(Cl-)=Ksp. If the
molar concentrations are taken as good approximations to the activities, which in dilute solutions
they are, then Ksp=[Ag+][Cl].
Example. Let us write and simplify to the extent possible the equilibrium constant for the equilibrium
Al3+(aq) + 3OH-(aq) <--> A1(OH) 3 (s) For this equilibrium K= 1/[Al3+][OH-]^3 = 1/Ksp. where K has the
units dm^12 /mol^4 , or (dm^3 )^4 /mol^4.
The form of equilibrium constant indicated as Ksp is called the solubility product constant or, more
commonly, the solubility product. This constant therefore must refer to the process of a solid going
into solution (solubility) rather than the reverse, precipitation of solid from solution. As a
consequence, the ions are products and appear in the numerator.
The value of the solubility product is temperature-dependent and is generally found to increase with
increasing temperature. As a consequence, the molar solubility of ionic salts generally increases
with increasing temperature. The extent of this increase varies from one salt to another.