Transcript Chemistry 1011

Chapter 17
Solubility Equilibria
Dr. Peter Warburton
[email protected]
http://www.chem.mun.ca/zcourses/1011.php
Solubility Equilibria
Dissolution
Precipitation
MmXx (s)  m Mn+ (aq) + x Xy- (aq)
m Mn+ (aq) + x Xy- (aq)  MmXx (s)
For a dissolution process, we give the
equilibrium constant expression the name
solubility product (constant) Ksp. For
MmXx (s)  m Mn+ (aq) + x Xy- (aq)
Ksp = [Mn+]m [Xy-]x
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Problem
Write the expressions for the
solubility product constant Ksp
of:
a) AgCl
b) PbI2
c) Ca3(PO4)2
d) Cr(OH)3
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Measuring Ksp and Calculating Molar
Solubility from Ksp
For the dissolution of solid CaF2 in water, we
might find that the concentrations of the ions
Ca2+ and F- at equilibrium are
[Ca2+] = 3.3 x 10-4 mol/L
[F-] = 6.7 x 10-4 mol/L
Ksp = [Ca2+] [F-]2
= (3.3 x 10-4 mol/L) (6.7 x 10-4 mol/L)2
= 1.5 x 10-10
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Measuring Ksp and Calculating Molar
Solubility from Ksp
We could also find the Ksp of a
solid by mixing solutions of known
concentrations of the ions, leading
to the precipitation of the solid.
When the system reaches
equilibrium we can measure the
ion concentrations to calculate
Ksp.
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The advantage of using the
dissolution of the solid
method is that a clear
relationship between the
concentrations of the ions
exists, based upon their relative
stoichiometry in the solid.
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If we mix two separate solutions that are
sources of calcium ions and fluoride ions, no
such relationship exists, and we can have
infinitely many mixtures that still obey the
solubility product expression.
Since Ksp values are equilibrium
constants, they will change with
temperature!
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Molar solubility and saturation
If we know the Ksp value for a solid, we
can calculate its molar solubility, which is
the number of moles of the solid that
can dissolve in the solvent before the
solution becomes saturated (no more
solid will dissolve).
Saturated means
equilibrium!
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Problem
A saturated solution of Ca3(PO4)2 has
[Ca2+] = 2.01 x 10-8 mol/L and
[PO43-] = 1.6 x 10-5 mol/L.
Calculate Ksp for Ca3(PO4)2
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Problem
If a saturated solution of BaSO4 is
prepared by dissolving solid BaSO4 in
water, and [Ba2+] = 1.05 x 10-5 mol/L,
what is the Ksp for BaSO4?
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Problem
Which has the greater molar
solubility?
AgCl with Ksp = 1.8 x 10-10
or
Ag2CrO4 with Ksp = 1.1 x 10-12
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Factors that Affect Solubility
The Common-Ion Effect
If a solution already has a
significant concentration of a
common-ion the solution will
dissolve LESS of the solid
than the same volume of pure
water can.
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The Common-Ion Effect
MmXx (s)  m Mn+ (aq) + x Xy- (aq)
Ksp = [Mn+]m [Xy-]x
Imagine that instead of dissolving
a solid in a solution of commonion that we add a common ion to
a solution created by dissolving
the solid in pure water.
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The Common-Ion Effect
MmXx (s)  m Mn+ (aq) + x Xy- (aq)
Ksp = [Mn+]m [Xy-]x
From Le Chatalier’s Principle, we know if we add
one of the product ions (the common ion we
added), the stress on the equilibrium is this
added product. To relieve the stress the
equilibrium will shift towards the reactants
meaning more solid is created.
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Problem
Calculate the molar
solubility of MgF2
-11
(Ksp = 7.4 x 10 )
in pure water and
in 0.10 mol/L MgCl2.
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Precipitation of Ionic Compounds
Can we predict if a solid will precipitate if
we mix two solutions of different ions?
YES!
Consider the mixing of a solution of Ca2+
ions and a solution of F- ions. A
precipitation of solid is the dissolution
reaction in reverse, so we can express
the reaction as
CaF2 (s)  Ca2+ (aq) + 2 F- (aq)
Ksp = [Ca2+ ][F-]2
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Precipitation of Ionic Compounds
When we mix the solutions, the system
is most likely not at equilibrium.
For solid dissolution / precipitation
reactions, we use a procedure similar to
the reaction quotient Qc to define the
ion product (IP) or Qsp.
Qsp = [Ca2+ ][F-]2
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If Qsp = Ksp, the solution is saturated, and the
system is at equilibrium.
If Qsp > Ksp, the solution is supersaturated, so
the system is not at equilibrium. The
concentration of the ions is greater than it
would be at equilibrium, and so the reaction
proceeds from ions towards the solid.
We expect precipitation to
occur!
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If Qsp < Ksp, the solution is
unsaturated, so the system is not at
equilibrium. The concentration of the
ions is less than it would be at
equilibrium, and so
we expect more solid to be able to
dissolve in this solution!
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Will a precipitate form when 0.150 L
of 0.10 molL-1 Pb(NO3)2 and 0.100 L
of 0.20 molL-1 NaCl are mixed?
Ksp of PbCl2 is 1.2 x 10-5
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Problem
Will a precipitate form on mixing equal
volumes of the following solutions?
a) 3.0 x 10-3 mol/L
BaCl2 and 2.0 x 10-3
mol/L Na2CO3
b) 1.0 x 10-5 mol/L Ba(NO3)2 and
4.0 x 10-5 mol/L Na2CO3
(Ksp of BaCO3 is 2.6 x 10-9)
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