Introductory University Chemistry I
Solutions and Solubility
Solubility 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 KClO4, which is
moderately soluble but is still sometimes used to precipitate potassium ion from
aqueous solutions. Metal nitrates are 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
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) <--> Al(OH)3(s)
For this equilibrium K = 1/[Al3+] [OH-]3 = 1/Ksp.
where K has the units dm12/mol4, or (dm3)4/mol4.
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. It is
sometimes possible to take advantage of the difference in the effect of
temperature to separate mixtures of different soluble salts. As the chart in
the following Figure shows, a solution originally of equal concentration in KClO3 and
KNO3 should upon heating and evaporation of water precipitate KClO3
because KNO3 is by far the more soluble near the boiling point of water.
Figure is not available.
The solubility of solid salts in water, and in most other solvents, increases
with temperature while that of gases decreases. This is an application of Le
Chatelier's principle, discussed in a preceding section. The heat or enthalpy
change of the dissolution reaction for most solids is positive so the dissolution
reaction is endothermic. For some solids, such as NaCl, the heat of solution is
very small and so the effect of temperature is small also. For other salts, such
as KNO3, the effect of temperature is much larger:
NaCl(c) <--> Na+(aq) + Cl-(aq); H0 = (-240.12 - 167.159) - (-411.153) = +3.87 kJ/mol
KNO3(c) <--> K+(aq) + NO3-(aq); H0 = (-252.38 - 205.0)-(-494.63) = +37.3 kJ/mol
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Copyright 1995 James A. Plambeck (Jim.Plambeck@ualberta.ca).
Updated November 6, 1995 jp.