•Solubility is a very important phenomenon.
•We can flavor foods due to the solubility of salt and
sugar in water.
•In this section, we will consider sparingly soluble
substances and quantitative ways to determine “How
soluble?”.
•According to the solubility rules previously learned,
PbCl2 and AgCl are both insoluble salts.
•However, if chloride ion is added to a solution containing both
Pb2+ and Ag+ ions, nearly all of the Ag+ is precipitated as AgCl
before any Pb2+ separates from the solution as PbCl2.
•This occurs because AgCl is much less soluble than PbCl2.
•To explain these differences in solubility, solubility equilibrium
must be examined quantitatively.
•When placed in water, a small amount of AgCl
dissolves and the following equilibrium is established
once the solution becomes saturated:
AgCl(s) ⇌Ag+(aq) + Cl-(aq)
Equilibrium expression:
If reactants or products have a coefficient other than
+
Ksp = [Ag ][Cl ] 1 concentrations must be raised to that power.
Ksp = solubility product constant
Ksp equals the product of concentration terms for the
ions dissolved in a saturated solution of a sparingly
soluble substance.
•The solubilities of salts change with temperature so a
value of Ksp applies only to solutions only at the
temperature at which its value was determined.
• Ksp can be obtained from a salt’s molar solubility in
water – the number of moles of solute dissolved in
one liter of its saturated solution.
• Examples
• Molar solubility can also be computed (estimated)
from values of Ksp.
• Examples
•Suppose we stir some calcium carbonate in water long
enough to establish the following equilibrium:
CaCO3 (s) ⇌ Ca2+ (aq) + CO32- (aq)
•Then we add to the solution a very soluble salt of
calcium, like CaCl2.
•This puts Ca2+ into solution, and it upsets the above
equilibrium.
•The ion product is no longer equal to Ksp.
•Remember from Le Chatelier’s principle, if we add
then equilibrium shifts opposite.
•The above equilibrium shifts to the left causing CO32to precipitate as CaCO3.
•Eventually equilibrium is reestablished, but with a lower
concentration of CO32+ in solution.
•In this new system, there are two sources of Ca2+, the
added CaCl2 and the CaCO3 still in solution.
•Because Ca2+ is common to both sources, it is called a
common ion.
•The addition of the common ion lowers the solubility of
CaCO3; it is less soluble in the presence of CaCl2 (or any
other soluble calcium salt) than it is in pure water.
•This lowering of the solubility of an ionic compound by
the addition of a common ion is called the common ion
effect.
•The common ion effect can dramatically lower the
solubility of a salt.
•So far we have considered solids dissolving in
solutions.
•Now we will consider the reverse process – the
formation of a solid from a solution.
•We will use the ion product.
•Ion product (Q) is defined like expression for Ksp but
uses initial concentrations instead of equilibrium
concentrations.
•Precipitate will form
•No precipitate will form
•No precipitate will form
Ion product >Ksp (supersaturated)
Ion product = Ksp (saturated)
Ion product < Ksp (unsaturated)
•The pH of a solution can affect a salt’s solubility.
•Example:
Mg(OH)2(s) ⇌ Mg2+(aq) + 2OH-(aq)
Add OH(increase basicity)
Decrease
Solubility
of Mg(OH)2
Add H+
(increase acidity -removes
OH- by reaction to produce H2O)
Increases
solubility (lose OHand must replace)
•General rule – If the anion X- is an effective base (HX is a weak
acid) the salt MX will show increased solubility in an acidic
solution.
•A complex ion is a
charged species consisting
of a metal ion surrounded
by ligands.
•A ligand is simply a Lewis
base.
•Recall a Lewis base is an ion or molecule having a
lone electron pair that can be donated to an empty
orbital on the metal ion to form a covalent bond.
•Some common ligands are H2O, NH3, Cl-, and CN-.
•Metal ions add ligands one at a time in steps
characterized by equilibrium constants called
formation constants or stability constants.
•For example, when solutions containing Ag+ and NH3
molecules are mixed, the following reactions take
place:
Ag+ (aq) + NH3 (aq) ⇌ Ag(NH3)+ (aq)
K1 = 2.1 x 103
Ag(NH3)+(aq) + NH3(aq)⇌ Ag(NH3)2+(aq)K2 = 8.2 x 103
where K1 and K2 are the formation constants for the
two steps.
•In a solution containing Ag+ and NH3, all the species
NH3, Ag+, Ag(NH3)+, and Ag(NH3)+ exist at equilibrium.
•When we write the formula for a complex, we follow
two rules:
1. The symbol for the metal ion is always given first,
followed by the ligands.
2. The charge on the complex is the algebraic sum of
the charge on the metal ion and the charges on the
ligands.
•
•
For example, the formula of the complex ion of Cu2+
and H2O is written Cu(H2O)42+ with the Cu first
followed by the ligands.
The charge on the complex is 2+ because the
copper ion has a charge of 2+ and the water
molecules are neutral.
•Metal ions that commonly form complex ions (or
coordination compounds): Al3+, Cu2+, Zn2+, Fe2+ (or
3+), Ni2+, Ag+. (All Curiously Colored Zebras Felt
Nicely Agreeable)
•Common ligands: NH3, OH-, Cl-, SCN-, CN-, H2O.
•Most common coordination number: twice the charge
of the metal ion.
•That means that Ag+ and NH3 → Ag(NH3)2+.
•Formation of complex ions is a common means to
dissolve otherwise insoluble salts.
•For example, in a solution with only water present,
AgCl only dissociates (dissolves) slightly to form Ag+.
•When ammonia is added, the Ag+ complexes with
the ammonia, and the removal of the Ag+ from the
solution as it converts to Ag(NH3)2+ pulls the AgCl
dissociation equilibrium to the right (LeChatelier’s
Principle).
•If sufficient ammonia is added to complex all of the
silver ions, the AgCl will completely dissolve.
AgCl(s) ⇌ Ag+ (aq) + Cl- (aq)