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Lewis Acids and Bases
There are reactions in nonaqueous solvents, in the gaseous state, and
even in the solid state that can be considered acid–base reactions in
which Brønsted–Lowry theory is not adequate to explain.
Why Lewis “acid”
and Lewis “base”?
Electron pair
acceptor
H+
+
Electron pair
donor
→
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The Solubility Product Constant, Ksp
‫מכפלת המסיסות‬
• Many important ionic compounds are only slightly soluble
(“insoluble”) in water.
• The equilibrium between the compound and the ions
present in a saturated aqueous solution can be written as
follows:
BaSO4(s)
Ba2+(aq) + SO42–(aq)
• Solubility product constant, Ksp: the equilibrium constant
expression for the dissolution of a slightly soluble solid.
Ksp = [ Ba2+ ][ SO42–]
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Ksp and Molar Solubility
• Ksp is an equilibrium constant
• Molar solubility is the number of moles of compound that
will dissolve per liter of solution.
‫מסיסות מולארית‬
• Molar solubility is related to the value of Ksp, but molar
solubility and Ksp are not the same thing.
– “smaller Ksp” doesn’t always mean “lower molar solubility”.
¾ The reason: different Ksp may have different units (Mx)
• Solubility depends on both Ksp and the form of the
equilibrium constant expression.
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The Common Ion Effect
in Solubility Equilibria
• The common ion effect affects solubility equilibria as it
does other aqueous equilibria.
• The solubility of a slightly soluble ionic compound is
lowered when a second solute that furnishes a common ion
is added to the solution.
Ag2SO4(s)
If MgSO4 is added to solution
⇒ Solubility of Ag2SO4 ↓
2 Ag+(aq) + SO4–2(aq)
VideoClip: AgI + NaI
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Common Ion Effect Illustrated
1. When Na2SO4(aq)
is added to the
saturated solution of
Ag2SO4 …
2. … [Ag+] attains a
new, lower equilibrium
concentration as Ag+
reacts with SO42– to
produce Ag2SO4.
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Does Precipitation Occur?
Qip is the ion product reaction quotient
and is based on initial conditions of the
reaction
Precipitation should occur if Qip > Ksp
Precipitation cannot occur if Qip < Ksp
A solution is just saturated if Qip = Ksp
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Is Precipitation Complete?
• We consider precipitation to be essentially
complete if about 99.9% of the target ion is
precipitated (<0.1% is left as free ions in solution)
• Conditions that generally favor completeness of
precipitation:
– A very small value of Ksp ⇒ few ions in solution
– A high initial concentration of the target ion
– A concentration of common ion that greatly exceeds
that of the target ion
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Selective Precipitation
A 2.00 M AgNO3 aqueous solution is slowly added from a buret to an
aqueous solution of 0.0100 M Cl– and 0.0100 M I–.
a. Which ion, Cl– or I–, is the first to precipitate from solution?
b. When the second ion begins to precipitate, what is the remaining
concentration of the first ion?
c. Is separation of the two ions by selective precipitation feasible?
AgCl(s)
Ag+(aq) + Cl–(aq)
Ksp = 1.8 x 10–10 M2
AgI(s)
Ag+(aq) + I–(aq)
Ksp = 8.5 x 10–17 M2
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Effect of pH on Solubility
• If the anion of a precipitate is that of a weak acid, the
precipitate will dissolve somewhat when the pH is lowered:
CaF2(s)
Ca2+(aq) + 2 F–(aq)
Added H+ will reacts with, and remove, F–;
LeChâtelier’s principle ⇒ more F– will form.
• If, however, the anion of the precipitate is that of a strong acid,
lowering the pH will have no effect on the precipitate.
AgCl(s)
Ag+(aq) + Cl–(aq)
Mg(OH)2 + HCl clip
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H+ will not consume Cl– ;
acid will not affect the
equilibrium.
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Equilibria Involving Complex Ions
Silver chloride becomes more
soluble, not less soluble, in high
concentrations of chloride ion.
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Complex Ion Formation
• A complex ion consists of a central metal atom or ion, with
other groups called ligands bonded to it.
– The metal ion accepts electron pairs (Lewis acid).
– Ligands act donate electron pairs (Lewis base).
– Common ligands:
• The formation of a complex ion is reversible.
• The equilibrium involving a complex ion, the metal ion, and
the ligands is described through a formation constant, Kf:
Ag+(aq) + 2 Cl–(aq)
[AgCl2]–(aq)
[AgCl2]–
Kf = –––––––––– = 1.2 x 108 M-2
[Ag+][Cl–]2
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Complex Ion Formation
1. Concentrated NH3
added to a solution of
pale-blue Cu2+ …
2. … forms deepblue Cu(NH3)42+.
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Some Formation Constants
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Complex Ion Formation
and Solubilities
2. But if the concentration
of NH3 is made high
enough …
3. … the AgCl forms the
soluble [Ag(NH3)2]+ ion.
1. AgCl is insoluble
in water.
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Complex Ions in Acid–Base Reactions
• Water molecules are commonly found as ligands in
complex ions (H2O is a Lewis base).
[Na(H2O)4]+
[Al(H2O)6]3+
[Fe(H2O)6]3+
• The electron-withdrawing power of a small, highly
charged metal ion can weaken an O—H bond in one of
the ligand water molecules.
• The weakened O—H bond can then give up its proton to
another water molecule in the solution.
• The complex ion acts as an acid.
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Ionization of a Complex Ion
The highly-charged iron(III)
ion withdraws electron density
from the O—H bonds.
[Fe(H2O)6]3+ + H2O
[Fe(H2O)5OH]2+ + H3O+
Ka = 1 x 10–7 M
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Amphoteric Species
• Certain metal hydroxides, insoluble in water, are amphoteric; they will
react with both strong acids and strong bases.
• Al(OH)3, Zn(OH)2, and Cr(OH)3 are amphoteric.
• Certain oxides are also amphoteric: Al2O3, ZnO, and Cr2O3.
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Summary of Concepts
• Lewis acid: accepts an electron pair
Lewis base: donates an electron pair
• The solubility product constant, Ksp, describes the equilibrium between
a slightly soluble ionic compound and its ions in a saturated aqueous
solution
• Precipitation is assumed to be complete if no more than 0.1% of the
target ion remains in solution
• Solubility is influenced by:
– The common ion effect (lowers solubility)
– pH (in some slightly soluble compounds)
– Complex ion formation (in some compounds)
• The ability of water as ligand molecules to donate protons accounts for
the acidic character of some complex ions
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