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Chapter Outline
• 15.1 Acids and Bases: The BrØnsted–Lowry Model
• 15.2 Acid Strength and Molecular Structure
• 15.3 pH and the Autoionization of Water
• 15.4 Calculations Involving pH, Ka, and Kb
• 15.5 Polyprotic Acids
• 15.6 pH of Salt Solutions
• 15.7 The Common-Ion Effect
• 15.8 pH Buffers
• 15.9 pH Indicators and Acid–Base Titrations
• 15.10 Solubility Equilibria
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pH of Salt Solutions
1. Neutral Salts (pH = 7) are from strong electrolytes (100%
ionization)
(a)Ionic Compounds:
NaCl(aq)  Na+(aq) + Cl-(aq)
base
conj. acid
HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
acid
conj. base
Infinitely
strong
Infinitely
weak
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pH of Salt Solutions
2. Basic Salts (pH > 7) are conjugate bases of weak acids
HClO(aq) + H2O(l) = H3O+(aq) + ClO-(aq)
conj.
weak
base
acid
ClO-(aq) + H2O(l) = OH-(aq) + HClO(aq)
conj. base
Ka (HClO) = 2.9 x 10-8
pH > 7
Kb = Kw
Ka
-14
Kb = 1.0 x 10
2.9 x 10-8
Kb (ClO-) = 3.4 x 10-7
pH of Salt Solutions
3. Acidic Salts (pH < 7) are conjugate acids of weak bases
NH3(aq) + H2O(l) = OH-(aq) + NH4+(aq)
conj.
weak
acid
base
NH4+(aq) + H2O(l) = H3O+(aq) + NH3(aq)
conj. acid
Kb (NH3) = 1.8 x 10-5
pH < 7
Ka = Kw
Kb
-14
Ka = 1.0 x 10
1.8 x 10-5
Ka (NH4+) = 5.6 x 10-10
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Calculating the pH of Solutions of Weak Acids and Bases:
Use the RICE Table as Before
1. Calculating the pH of a Solution of a Basic Salt
ClO-(aq) + H2O(l) = OH-(aq) + HClO(aq)
0.100 M - x
x
Kb = 3.4 x 10-7 =
3.4 x 10-7 =
x =
x
x2
0.100 M - x
x2
0.100 M
Ka = 2.9 x 10-8
Kb = Kw
Ka
Kb = 3.4 x 10-7
Since Ka is < 10-5, assume
that x << 0.100 M
(3.4 x 10-7)(0.100)
x = [OH-] = 1.9 x 10-4 M
pH = - log (1.9 x 10-4) = 3.7
Assumption OK
pH = 14 - 3.7 = 10.3
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Calculating the pH of Solutions of Weak Acids and Bases:
Use the RICE Table as Before
2. Calculating the pH of a Solution of an Acidic Salt (Ex. 15.8)
What is the pH of a 0.25 M solution of NH4Cl?
Chapter Outline
• 15.1 Acids and Bases: The BrØnsted–Lowry Model
• 15.2 Acid Strength and Molecular Structure
• 15.3 pH and the Autoionization of Water
• 15.4 Calculations Involving pH, Ka, and Kb
• 15.5 Polyprotic Acids
• 15.6 pH of Salt Solutions
• 15.7 The Common-Ion Effect
• 15.8 pH Buffers
• 15.9 pH Indicators and Acid–Base Titrations
• 15.10 Solubility Equilibria
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The Common Ion Effect
A shift in equilibrium caused by the addition of a compound
having an ion in common with the dissolved substance.
CH3COONa (s)
Na+ (aq) + CH3COO- (aq)
CH3COOH (aq)
H+ (aq) + CH3COO- (aq)
common
ion
The common ion effect can be used to produce a BUFFER
SOLUTION = a solution of a weak acid or base and it's
conjugate, e.g. CH3COOH and CH3COONa
By controlling the ratio of weak acid/base to it's conjugate,
we can shift the equilibrium to whatever [H+] and therefore
pH we want.
The Henderson-Hasselbach Equation
HA(aq) + H2O(aq) = H3O+(aq) + A-(aq)
Ka =
[H3O+][A-]
[HA]
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1. Solution using the RICE table
2. Solution using the Henderson-Hasselbach equation
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