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General Chemistry II
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SOLUBILITY AND PRECIPITATION
EQUILIBRIA
16.1 The Nature of Solubility Equilibria
16.2 Ionic Equilibria between Solids and Solutions
16.3 Precipitation and the Solubility Product
16.4 The Effects of pH on Solubility
16.5 Complex Ions and Solubility
16.6 Selective Precipitation of Ions
16CHAPTER
General Chemistry II
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General Features of Solubility Equilibria
Saturation
~ Dissolution-precipitation equilibrium
Fig. 16.1 Deposit of K2PtCl4 from the saturated
aqueous solution as the water evaporates.
16.1 THE NATURE OF SOLUBILITY EQUILIBRIA
Recrystallization ~ Purification of solids
Solvent of crystallization
2 Li+(aq) + SO42-(aq) + H2O(l) → Li2SO4H2O(s)
~ different chemical formula & mass
Supersaturation ~ Slow equilibrium
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The solubility of Ionic Solids
Fig. 16.3 Temperature dependence of solubility.
Solubility at 25°C,
AgClO4 ; 5570 g/L, AgCl; 0.0018 g/L
Temperature dependence
- Mostly endothermic
→ Solubility increases with T
- CaSO4 exothermic
→ Solubility decreases with T
Classification (at 25 °C)
Soluble > 10 g/L,
Slightly soluble 0.1~10 g/L,
Insoluble < 0.1 g/L
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Highly soluble salt: Nonideal solution, CsCl(s) Cs+(aq) + Cl-(aq)
Fig. 16.5 The dissolution of the ionic solid CsCl in water
16.2 IONIC EQUILIBRIA BETWEEN SOLIDS AND
SOLUTIONS
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Solubility and Ksp
Solubility product:
Ksp = [Ag+][Cl-] = 1.610-10 at 25 °C
AgCl(s) Ag+(aq) + Cl-(aq)
Solubility (S) of AgCl at 25°C calculated from Ksp
Ksp = [Ag+][Cl-] = S2 = 1.610-10
S = 1.2610-5 M
Gram solubility = (1.2610-5 mol/L) (143.3 g/mol)
= 1.810-3 g/L
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CaF2(s) Ca2+(aq) + 2 F-(aq)
Ksp = [Ca2+][F-]2 = 3.910-11 at 25°C
[Ca2+] = S, [F-] = 2S
Ksp = [Ca2+][F-]2 = S (2S)2 = 4S3 → S = 2.1 10-4 M
Gram solubility = (2.110-4 mol/L) (78.1 g/mol) = 0.017 g/L
EXAMPLE 16.1 Calculation of [Ca2+] and [F-] in a saturated solution
of CaF2 at 25°C: Ksp → Solubility
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Ksp = [Ag+]2[CrO42-] = 2.7 10-12
Gram solubility: 0.029 g/L
Molar solubility: 0.029 g/L = 8.74 10-5 mol/L = S
[Ag+]= 2S, [CrO42-] = S
Ksp = [Ag+]2[CrO42-] = 4S3 = 2.7 10-12
→ 42 % greater than the tabulated value, 1.9 10-12
Solubility (0.029 g/L) → Ksp
Ag2CrO4(s) 2 Ag+(aq) + CrO42-(aq)
EXAMPLE 16.2
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Fig. 16.6 A plot of precipitation and dissolution equilibrium for AgCl in water.
The slope of the path toward equilibrium represented by red or blue arrow is 1.
16.3 PRECIPITATION AND THE SOLUBILITY
PRODUCT
Precipitation from Solution
Ksp = [Ag+][Cl-]
Q0 = [Ag+]0[Cl-]0
~ initial reaction quotient
Q0 > Ksp precipitation
Q0 < Ksp dissolution
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Cl- is the limiting reactant → complete precipitation first
Remaining [Ag+] = 0.0015 - 5.0 10-6 0.0015 M
AgCl(s) Ag+(aq) + Cl-(aq)----------------------------------------------------------------------Initial 0.0015 0Change + y + y
--------------- ------Equilibrium 0.0015 + y y----------------------------------------------------------------------
Ksp = 1.60 10-10 = (0.0015 + y) y 0.0015 y
y = [Cl-] = 1.1 10-7 M, [Ag+] = 0.0015 M
[Ag+]0 = 0.0015 M, [Cl-]0 = 5.010-6 M
Equilibrium concentrations?
EXAMPLE 16.4
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The Common-Ion Effect
~ Solubility decreases in the presence of a common ion
AgCl NaCl or AgNO3
EX. Solubility of AgCl(s) in 1.00 L of 0.100 M NaCl solution
[Ag+]NaCl = S, [Cl-]NaCl = 0.100 + S
Ksp = 1.60 10-10
= [Ag+] NaCl [Cl-] NaCl
= S (0.100 + S) 0.100 S
(S < Swater =1.3 10-5 << 0.100)
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Fig. 16.7 Common-ion effect for the solubility
of AgCl in AgNO3 solution and in NaCl solution.
2
5H O 3
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0.1M NaCl
[Ag ] 1.3 108.1 10
[Ag ] 1.6 10
[Ag+] NaCl = S = 1.60 10-9 M
[Cl-] NaCl = 0.100 M
2
5
H O[Ag ] 1.3 10 M
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Fig. 16.8 Damage due to increased acidity from air pollution.
On the east pier of Stanford White's Washington Square Arch
is Herma A. MacNeil's Washington in War (1916)
(Washington Square Park in the Greenwich Village neighborhood
of Lower Manhattan in New York City)
CaCO3(s) + H3O+(aq) Ca2+(aq) + HCO3
-(aq) + H2O(l)
16.4 THE EFFECTS OF pH ON SOLUBILITY
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Solubility of Hydroxides
In pure water, [Zn2+] = S, [OH-] = 2S Ksp = S(2S)2
S = [Zn2+] = 2.2 10-6 M, [OH-] = 2S = 4.5 10-6 M, pH = 8.65
EXAMPLE 16.6
Zn(OH)2(s) Zn2+(aq) + 2 OH-(aq)
Ksp = [Zn2+][OH-]2 = 4.5 10-17
In acidic solution, [OH-] decreases. → reaction goes to the right
Comparison of solubilities of Zn(OH)2(s) in pure water
and in a buffer with pH 6.00.
In a pH = 6.00 buffer, [OH-] = 1.0 10-8 M (fixed).
[Zn2+] = Ksp / [OH-]2 = 0.45 M
Metal hydroxides are basic → more soluble in acidic solution
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Solubility of Salts of Bases
CaF2(s) Ca2+(aq) + 2 F-(aq), Ksp = 3.9 10-11
- Solubility of CaF2(s) at low pH :
F-(aq) + H3O+(aq) HF(aq) + H2O(l), K = 2.9 103
more soluble in acidic solution (large K)
[H3O+] → [F-] → more CaF2(s) dissolves (Le Chatelier)
- Solubility of AgCl(s) at low pH :
AgCl(s) Ag+(aq) + Cl-(aq)
- Even in acidic solution,
Cl-(aq) + H3O+(aq) HCl(aq) + H2O(l)
→ negligible effect of pH on the solubility of AgCl
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Problem Sets
For Chapter 16,
14, 22, 30, 34
Recommended