2/2/10
Acids
Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon
dioxide gas
Acids and Bases
Bases
Chapter 15
Have a bitter taste.
Feel slippery. Many soaps contain bases.
4.3
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Arrhenius acid is a substance that produces H+ (H3O+) in water
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
Arrhenius base is a substance that produces OH- in water
base
acid
acid
base
base
acid
conjugate
acid
conjugate
base
4.3
15.1
The Ion Product of Water
Acid-Base Properties of Water
H+ (aq) + OH- (aq)
H2O (l)
H2O (l)
autoionization of water
H
O
H
+ H
[H
O
H
]
H
+
H2O + H2O
Kc =
[H+][OH-]
[H2O]
[H2O] = constant
Kc[H2O] = Kw = [H+][OH-]
+ H
O
-
The ion-product constant (Kw) is the product of the molar
concentrations of H+ and OH- ions at a particular temperature.
H
base
acid
O
H+ (aq) + OH- (aq)
Solution Is
conjugate
acid
H3O+ + OHconjugate
base
At 250C
Kw = [H+][OH-] = 1.0 x 10-14
15.2
[H+] = [OH-]
neutral
[H+] > [OH-]
acidic
[H+] < [OH-]
basic
15.2
1
2/2/10
pH – A Measure of Acidity
What is the concentration of OH- ions in a HCl solution
whose hydrogen ion concentration is 1.3 M?
pH = -log [H+]
Kw = [H+][OH-] = 1.0 x 10-14
Solution Is
[H+] = 1.3 M
[OH-] =
Kw
1 x 10-14
=
= 7.7 x 10-15 M
[H+]
1.3
At 250C
neutral
[H+] = [OH-]
[H+] = 1 x 10-7
pH = 7
acidic
[H+] > [OH-]
[H+] > 1 x 10-7
pH < 7
basic
[H+] < [OH-]
[H+] < 1 x 10-7
pH > 7
pH
[H+]
15.2
15.3
The pH of rainwater collected in a certain region of the
northeastern United States on a particular day was 4.82.
What is the H+ ion concentration of the rainwater?
pH = -log [H+]
pOH = -log [OH-]
[H+]
[H+][OH-] = Kw = 1.0 x 10-14
-log [H+] – log [OH-] = 14.00
= 10-pH = 10-4.82 = 1.5 x 10-5 M
The OH- ion concentration of a blood sample is 2.5 x 10-7 M.
What is the pH of the blood?
pH + pOH = 14.00
pH + pOH = 14.00
pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60
pH = 14.00 – pOH = 14.00 – 6.60 = 7.40
15.3
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Weak Acids are weak electrolytes
Na+ (aq) + Cl- (aq)
HF (aq) + H2O (l)
Weak Electrolyte – not completely dissociated
CH3COOH
CH3
COO-
(aq) +
H+
(aq)
HNO3 (aq) + H2O (l)
H3
(aq)
H3O+ (aq) + NO2- (aq)
HSO4- (aq) + H2O (l)
H3O+ (aq) + SO42- (aq)
H3O+ (aq) + OH- (aq)
Strong Bases are strong electrolytes
NaOH (s)
H3O+ (aq) + Cl- (aq)
O+
H3O+ (aq) + F- (aq)
HNO2 (aq) + H2O (l)
H2O (l) + H2O (l)
Strong Acids are strong electrolytes
HCl (aq) + H2O (l)
15.3
KOH (s)
+ NO3 (aq)
-
HClO4 (aq) + H2O (l)
H3O+ (aq) + ClO4- (aq)
H2SO4 (aq) + H2O (l)
H3O+ (aq) + HSO4- (aq)
H 2O
H 2O
Ba(OH)2 (s)
15.4
Na+ (aq) + OH- (aq)
K+ (aq) + OH- (aq)
H 2O
Ba2+ (aq) + 2OH- (aq)
15.4
2
2/2/10
Weak Bases are weak electrolytes
F- (aq) + H2O (l)
OH- (aq) + HF (aq)
NO2 (aq) + H2O (l)
-
OH- (aq) + HNO2 (aq)
Conjugate acid-base pairs:
• 
The conjugate base of a strong acid has no measurable
strength.
• 
H3O+ is the strongest acid that can exist in aqueous
solution.
• 
The OH- ion is the strongest base that can exist in aqeous
solution.
15.4
Strong Acid
Weak Acid
15.4
What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
Start 0.002 M
HNO3 (aq) + H2O (l)
End 0.0 M
0.0 M
0.0 M
H3O+ (aq) + NO3- (aq)
0.002 M 0.002 M
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Start 0.018 M
Ba(OH)2 (s)
End 0.0 M
0.0 M
0.0 M
Ba2+ (aq) + 2OH- (aq)
0.018 M 0.036 M
pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56
15.4
15.4
15.5
15.5
Weak Acids (HA) and Acid Ionization Constants
HA (aq) + H2O (l)
H3O+ (aq) + A- (aq)
H+ (aq) + A- (aq)
HA (aq)
Ka =
[H+][A-]
[HA]
Ka is the acid ionization constant
Ka
weak acid
strength
3
2/2/10
What is the pH of a 0.5 M HF solution (at 250C)?
[H+][F-]
= 7.1 x 10-4
Ka =
HF (aq)
H+ (aq) + F- (aq)
[HF]
Initial (M)
Change (M)
Equilibrium (M)
HF (aq)
H+ (aq) + F- (aq)
0.50
0.00
-x
+x
+x
x
x
x2
= 7.1 x 10-4
Ka =
0.50 - x
Ka ≈
x2
= 7.1 x 10-4
0.50
[H+] = [F-] = 0.019 M
[HF] = 0.50 – x = 0.48 M
Ka << 1
x = 0.019
0.50 – x ≈ 0.50
x2 = 3.55 x 10-4
x = 0.019 M
pH = -log [H+] = 1.72
15.5
Solving weak acid ionization problems:
• 
0.50 – x ≈ 0.50
0.019 M
x 100% = 3.8%
0.50 M
Less than 5%
Approximation ok.
What is the pH of a 0.05 M HF solution (at 250C)?
x2
Ka ≈
= 7.1 x 10-4 x = 0.006 M
0.05
More than 5%
0.006 M
x 100% = 12%
0.05 M
Approximation not ok.
Must solve for x exactly using quadratic equation or method of
successive approximation.
15.5
What is the pH of a 0.122 M monoprotic acid whose
Ka is 5.7 x 10-4?
1.  Identify the major species that can affect the pH.
• 
Ka << 1
When x is less than 5% of the value from which it is subtracted.
0.00
0.50 - x
When can I use the approximation?
In most cases, you can ignore the autoionization of
water.
HA (aq)
H+ (aq) + A- (aq)
0.122
0.00
0.00
-x
+x
+x
0.122 - x
x
x
Initial (M)
Ignore [OH-] because it is determined by [H+].
2.  Use ICE to express the equilibrium concentrations in terms
of single unknown x.
3.  Write Ka in terms of equilibrium concentrations. Solve for x
by the approximation method. If approximation is not valid,
solve for x exactly.
4.  Calculate concentrations of all species and/or pH of the
solution.
Change (M)
Equilibrium (M)
Ka =
Ka ≈
x2
= 5.7 x 10-4
0.122 - x
x2
0.122
= 5.7 x 10-4
Ka << 1
0.122 – x ≈ 0.122
x2 = 6.95 x 10-5
x = 0.0083 M
More than 5%
0.0083 M
x 100% = 6.8%
0.122 M
Approximation not ok.
15.5
Ka =
x2
= 5.7 x 10-4
0.122 - x
ax2 + bx + c =0
x = 0.0081
Initial (M)
Change (M)
Equilibrium (M)
x2 + 0.00057x – 6.95 x 10-5 = 0
-b ± √b2 – 4ac
x=
2a
percent ionization =
H+ (aq) + A- (aq)
0.122
0.00
0.00
-x
+x
+x
0.122 - x
x
x
Ionized acid concentration at equilibrium
Initial concentration of acid
x 100%
For a monoprotic acid HA
x = - 0.0081
HA (aq)
[H+] = x = 0.0081 M
15.5
Percent ionization =
[H+]
[HA]0
x 100%
[HA]0 = initial concentration
pH = -log[H+] = 2.09
15.5
15.5
4
2/2/10
Weak Bases and Base Ionization Constants
NH3 (aq) + H2O (l)
NH4+ (aq) + OH- (aq)
Kb =
[NH4+][OH-]
[NH3]
Kb is the base ionization constant
weak base
strength
Kb
Solve weak base problems like weak acids
except solve for [OH-] instead of [H+].
15.6
15.6
Ionization Constants of Conjugate Acid-Base Pairs
HA (aq)
A- (aq)
H+ (aq) + A- (aq)
OH- (aq)
+ H2O (l)
Ka
+ HA (aq)
H+ (aq) + OH- (aq)
H2O (l)
Kb
Kw
KaKb = Kw
Weak Acid and Its Conjugate Base
Ka =
Kw
Kb
Kb =
Kw
Ka
15.7
15.8
Molecular Structure and Acid Strength
H X
H+ + X-
The
stronger
the bond
The
weaker
the acid
Molecular Structure and Acid Strength
Z
δO
δ+
H
Z
O- + H+
The O-H bond will be more polar and easier to break if:
• 
Z is very electronegative or
• 
Z is in a high oxidation state
HF << HCl < HBr < HI
15.9
15.9
5
2/2/10
Molecular Structure and Acid Strength
Molecular Structure and Acid Strength
1. Oxoacids having different central atoms (Z) that are from the
same group and that have the same oxidation number.
Acid strength increases as the oxidation number of Z increases.
••
••
••
••
••
••
Acid strength increases with increasing electronegativity of Z
••
••
O
O
••
••
••
••
H O Cl O
H O Br O
•• •• ••
•• •• ••
2. Oxoacids having the same central atom (Z) but different
numbers of attached groups.
Cl is more electronegative than Br
HClO4 > HClO3 > HClO2 > HClO
HClO3 > HBrO3
15.9
15.9
Acid-Base Properties of Salts
Acid-Base Properties of Salts
Neutral Solutions:
Acid Solutions:
Salts containing an alkali metal or alkaline earth metal
ion (except Be2+) and the conjugate base of a strong
acid (e.g. Cl-, Br-, and NO3-).
NaCl (s)
H 2O
Na+
(aq) +
Cl-
Salts derived from a strong acid and a weak base.
NH4Cl (s)
(aq)
NH4+ (aq)
Basic Solutions:
Salts derived from a strong base and a weak acid.
NaCH3COOH (s)
H 2O
CH3COO- (aq) + H2O (l)
H 2O
NH4+ (aq) + Cl- (aq)
NH3 (aq) + H+ (aq)
Salts with small, highly charged metal cations (e.g. Al3+,
Cr3+, and Be2+) and the conjugate base of a strong acid.
Na+ (aq) + CH3COO- (aq)
3+
Al(H2O)6 (aq)
CH3COOH (aq) + OH- (aq)
2+
Al(OH)(H2O)5 (aq) + H+ (aq)
15.10
15.10
Acid-Base Properties of Salts
Acid Hydrolysis of Al3+
Solutions in which both the cation and the anion hydrolyze:
15.10
• 
Kb for the anion > Ka for the cation, solution will be basic
• 
Kb for the anion < Ka for the cation, solution will be acidic
• 
Kb for the anion ≈ Ka for the cation, solution will be neutral
15.10
6
2/2/10
Definition of An Acid
Oxides of the Representative Elements
In Their Highest Oxidation States
Arrhenius acid is a substance that produces H+ (H3O+) in water
A Brønsted acid is a proton donor
A Lewis acid is a substance that can accept a pair of electrons
A Lewis base is a substance that can donate a pair of electrons
••
•• H+ + OH
••
acid base
H
15.11
acid
N H
••
H+ +
H
base
••
H O H
••
H
+
H N H
H
15.12
Lewis Acids and Bases
H
+
N H
••
F
F B
F
H
acid
base
F
F B
F
H
N H
H
No protons donated or accepted!
15.12
7