Transcript Slide 1

Chapter 9
Dissolution and Precipitation
Equilibria
9-1 The Nature of Solubility
Equilibria
9-2 The Solubility of Ionic Solids
9-3 Precipitation and the
Solubility Product
9-4 The Effects of pH on
Solubility
9-5 Complex Ions and Solubility
9-6 Controlling Solubility in
Qualitative Analysis
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Saturated Solution: a solution in
equilibrium with excess solute
e.g., NaCl solubility in grams
per 100 grams water is
approximately 36.0 grams =
saturated solution
Unsaturated Solution:
contains less than the
equilibrium concentration of
the solute
Supersaturated Solution: a
solution that temporarily
contains more of a solute than
the equilibrium quantity
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Endothermic – heat is added to a system
Exothermic – heat is removed from a system
Sharp
changes in
slope occur if
water of
crystallization
is lost or
gained by the
solid that is in
contact with
the solution
AgF has two
such changes
AgF·4H20
AgF·2H20
AgF
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9-2 Solubility of Salts
This chapter considers only salts
which are sparingly soluble or
insoluble for which concentrations of
saturated salts are [salt] = 0.1 Mol L-1
or less
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Solubility Product Ksp
Describes a chemical
equilibrium in which an excess
solid salt is in equilibrium with
a saturated aqueous solution of
its separated ions.
General equation
AB (s) ↔ A+ (aq) + B- (aq)
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The Solubility of Ionic Solids
The Solubility Product
AgCl(s) ↔Ag+ (aq) + Cl-(aq)
Ksp =
=
The solid AgCl, which is in excess, is
understood to have a concentration of
1 mole per liter.
Ksp= 1.6  10-10 at 25oC
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The Solubility of Ionic Solids
The Solubility Product
Ag2SO4(s) ↔2Ag+(aq) + SO42-(aq)
Ksp =
Fe(OH)3(s) ↔Fe+3(aq) + 3OH-1(aq)
Ksp =
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The Solubility of Ionic Solids
The Solubility Product
Exercise 9-1
Write the Ksp equation for the
dissolution of aluminum hydroxide
(Al(OH)3) in water.
Al(OH)3(s) ↔Al3+(aq) + 3 OH-(aq)
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The Solubility of Ionic Solids
The Solubility Product
TABLE 9-1contains
Ksp values at 25C
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The Solubility of Salts
Solubility and Ksp
Exercise 9-2
Determine the mass of lead(II) iodate dissolved in
2.50 L of a saturated aqueous solution of Pb(IO3)2 at
25oC. The Ksp of Pb(IO3)2 is 2.6  10-13.
Pb(IO3)2(s) ↔Pb2+(aq) + 2 IO3-(aq)
[y]
[Pb2+][IO3-]2 = Ksp
[y]
[2y]
y = 4.0  10-5 [Pb(IO3)2] = [Pb2+] = y = 4.0  10-5 mol L-1
 [IO3-] = 2y = 8.0  10-5 mol L-1
Gram solubility of = (4.0  10-5 mol L-1)  (557 g mol-1)
= 0.0223 g L-1  2.50 L
Lead (II) iodate
Pb=207.2
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I=126.9
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Pb(IO3)2 = 557g per mole
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The Solubility of Salts
Solubility and Ksp
Exercise 9-3
Compute the Ksp of silver sulfate (Ag2SO4) at 25oC if its
mass solubility is 8.3 g L-1.
1 Ag2SO4(s) ↔2 Ag+(aq) + 1 SO42-(aq)
[y]
[2y]
[y]
1. Given mass solubility
2. Find molar solubility
3. Find equilibrium expression
expressed in molar solubility
4. Find K sp
mass
moles 
Molar Mass
mass  moles x Molar Mass
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The Solubility of Salts
Solubility and Ksp
Exercise 9-3
Compute the Ksp of silver sulfate (Ag2SO4) at 25oC if its
mass solubility is 8.3 g L-1.
1 Ag2SO4(s) ↔2 Ag+(aq) + 1 SO42-(aq)
[y]
[2y]
[y]
[y] = (8.3 g Ag2SO4 L-1)  (1 mol Ag2SO4/311.8 g)
[y] = [2.66  10-2 ] mol Ag2SO4 L-1
[Ag+]2[SO42-] = Ksp
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Review
The Nature of Solubility Equilibria
Dissolution and precipitation are
reverse of each other.
Dissolution (Solubility)
General reaction
X3Y2 (s) ↔ 3X+2 (aq) + 2Y-3 (aq)
[s]
[3s]
Ksp = [X+2]3 [Y-3]2
[2s]
= (3s)3 (2s)2
s = molar solubility expressed in moles per liter
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Review
The Nature of Solubility Equilibria
Dissolution and precipitation are
reverse of each other.
Precipation
General reaction
X3Y2 (s) ↔ 3X+2 (aq) + 2Y-3 (aq)
Mix [X+2] and [Y-3] [X+2]
[Y-3]
Does a ppt of
X3Y2 form?
Reaction quotient before mixing occurs:
Q(init) = [X+2]3(init)[Y-3]2(init)
2 9
+2]3 [Y
-3Chapter
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K
[X
]
sp
Q>K?
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Precipitation from Solution: Does a solid ppt form?
AgCl(s) ↔Ag+ (aq) + Cl-(aq)
Q (init) =
[Ag+] (init) [Cl-] (init)= Reaction quotient
Ksp =
[Ag+][Cl-]
If Q > Ksp then the solid precipitates
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Precipitation and the Solubility
Product
Precipitation from Solution
Exercise 9-4:
The Ksp of thallium (I) iodate is 3.1  10-6 at 25oC.
Suppose that 555 mL of a 0.0022 M solution of TlNO3 is
mixed with 445 mL of a 0.0022 M solution of NaIO3.
Does TlIO3 precipitate at equilibrium?
Evaluate : Reaction quotient before mixing occurs:
Q(init) = [Tl+](init)[IO3-](init)
If Q(init) > Ksp, solid TlIO3 precipitates
until Q = Ksp
If Q(init) < Ksp, no solid TlIO3 can appear.
[Tl+]
Q > Ksp
Solid ppt
Q < Ksp
No ppt
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[IO3-]
Exercise 9-4
The Ksp of thallium(I) iodate is 3.1  10-6 at 25oC.
Suppose that 555 mL of a 0.0022 M solution of TlNO3 is
mixed with 445 mL of a 0.0022 M solution of NaIO3.
Does TlIO3 precipitate at equilibrium?
Tl(IO3) (s) ↔Tl+(aq) + IO3-(aq)
[Tl+](init) = (0.0022 mol L-1)(555 mL/1000 mL)
= 0.0012 mol L-1
[IO3-](init) = (0.0022 mol L-1)(445 mL/1000 mL)
= 0.00098 mol L-1
Q(init) = [Tl+](init)[IO3-](init)
= (0.0012)(0.00098) = 1.17  10-6
Q(init) ? Ksp
= 1.17  10-6 < 3.1  10-6
Because Q(init) < Ksp,
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solid TlIO3 does NOT
precipitate!
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Precipitation and the Solubility
Product
The Common Ion Effect
If a solution and a solid salt to
be dissolved in it have an ion in
common, then the solubility of
the salt is depressed.
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The Common Ion Effect
Exercise 9-6
The Ksp of thallium(I) iodate (TlO3) is 3.1
 10-6 at 25oC. Determine the molar
solubility of TlIO3 in 0.050 mol L-1 KIO3 at
25oC.
TlIO3(s) ↔ Tl+(aq) +
[Tl+] (mol L-1)
IO3-(aq)
[IO3-] (mol L-1)
Initial concentration
Change in concentration
Equilibrium concentration
[Tl+][IO3-] = Ksp
Assume s is small
s = [TlIO3]= 6.2 × 10-5 mol L-1 which is
-3
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depressed
28 times OFB
relative
conc. without the common ion
The Effects of pH on Solubility
Solubility of Hydroxides
Many solids dissolve more readily in
more acidic solutions
Zn(OH)2(s) ↔Zn2+(aq) + 2 OH-(aq)
[Zn2+][OH-]2 = Ksp = 4.5  10-17
If pH decreases (or made more acidic), the [OH-]
decreases. In order to maintain Ksp the [Zn2+]
must increase and consequently more solid
Zn(OH)2 dissolves.
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The Effects of pH on Solubility
Solubility of Hydroxides
Exercise 9-7
Estimate the molar solubility of Fe(OH)3 in a solution
that is buffered to a pH of 2.9.
[OH-] = 3y
In pure water: [Fe3+] = y
y(3y)3 = 27y4 = Ksp = 1.1  10-36
y = 4.5  10-10 mol L-1 = [Fe3+] = [Fe(OH)3]=
[OH-] = 3y = 1.3  10-9 mol L-1
pOH = 8.87
(and pH = 5.13)
In pure water, Fe(OH)3 is 5 x 10 6 less
soluble
than at pHOFB
= 2.9
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9-7 The Effects of pH on Solubility
•
•
•
•
Solubilities of Hydroxides
Solubility of Salts and Weak Bases
Selective Precipitation of Ions
Metal Sulfides
Somewhat more complicated due to other
competing reactions. E.g.,
MS + H2O ↔ M2+ + OH- + HS(Metal Sulfide)
But as before solubility of Metal
Sulfides increase as pH decreases
Ksp = [M2+][OH-][HS-]
As pH decreases (or made more acidic), the [OH-]
decreases. In order to maintain Ksp the [M2+] must
increase and consequently more solid Metal Sulfide
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dissolves.
• Examples / Exercises
– 9-1,
– 9-2,
– 9-3,
– 9-4,
– 9-5,
– 9-6,
– 9-7
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Ksp calculations
Ksp calculations
Ksp calculations
ppt Q ? Ksp
Equilibrium concentrations
Common Ion effect
Effect of pH of on solubility
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