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

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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.

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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