Chapter 9
Lecture
Outline
Prepared by
Andrea D. Leonard
University of Louisiana at Lafayette
1
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
9.1 Introduction to Acids and Bases
The Brønsted–Lowry definition:
•A Brønsted–Lowry acid is a proton (H+) donor.
•A Brønsted–Lowry base is a proton (H+) acceptor.
This proton is donated.
HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
•HCl is a Brønsted–Lowry acid because it donates
a proton to the solvent water.
•H2O is a Brønsted–Lowry base because it accepts
a proton from HCl.
4
9.1 Introduction to Acids and Bases
A. Brønsted–Lowry Acids
•A Brønsted–Lowry acid must contain a hydrogen atom.
•Common Brønsted–Lowry acids (HA):
HCl
hydrochloric acid
HBr
hydrobromic acid
HNO3
nitric acid
H2SO4
sulfuric acid
H
H
C
O
C
acidic H
atom
O H
H
acetic acid
5
9.1 Introduction to Acids and Bases
A. Brønsted–Lowry Acids
•A monoprotic acid contains one acidic proton.
HCl
•A diprotic acid contains two acidic protons.
H2SO4
•A triprotic acid contains three acidic protons.
H3PO4
•A Brønsted–Lowry acid may be neutral or it may
carry a net positive or negative charge.
HCl, H3O+, HSO4−
6
9.1 Introduction to Acids and Bases
B. Brønsted–Lowry Bases
•A Brønsted–Lowry base is a proton acceptor,
so it must be able to form a bond to a proton.
•A base must contain a lone pair of electrons that
can be used to form a new bond to the proton.
This e− pair forms a new
bond to a H from H2O.
H
N
H
+
H
Brønsted–Lowry
base
H2O(l)
+
H
H
N
H
+ OH− (aq)
H
7
9.1 Introduction to Acids and Bases
B. Brønsted–Lowry Bases
•Common Brønsted–Lowry Bases (B ):
Lone pairs make these
neutral compounds bases.
NH3
ammonia
H2O
water
The OH− is the base in
each metal salt.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Mg(OH)2
magnesium hydroxide
Ca(OH)2
calcium hydroxide
8
9.2 The Reaction of a Brønsted–Lowry
Acid with a Brønsted–Lowry Base
This e− pair
stays on A.
H A
acid
This e− pair forms
a new bond to H+.
gain of H+
+
B
base
A
−
+
H
B+
loss of H+
9
9.2 The Reaction of a Brønsted–Lowry
Acid with a Brønsted–Lowry Base
gain of H+
H A
acid
+
B
base
A − +
H B+
conjugate conjugate
acid
base
loss of H+
•The product formed by loss of a proton from an
acid is called its conjugate base.
•The product formed by gain of a proton by a base
is called its conjugate acid.
10
9.2 The Reaction of a Brønsted–Lowry
Acid with a Brønsted–Lowry Base
gain of H+
H Br
acid
+
H2O
base
Br− +
H3O+
conjugate conjugate
acid
base
loss of H+
•HBr and Br− are a conjugate acid–base pair.
•H2O and H3O+ are a conjugate acid–base pair.
•The net charge must be the same on both sides
of the equation.
11
9.2 The Reaction of a Brønsted–Lowry
Acid with a Brønsted–Lowry Base
Amphoteric compound: A compound that contains
both a hydrogen atom and a lone pair of e−; it can
be either an acid or a base.
+
H
H
O
H
add H+
H2O as a base
H
O
H
H2O as an acid
H
O
H
conjugate acid
remove
H+
−
H
O
conjugate base
13
9.3 Acid and Base Strength
A. Relating Acid and Base Strength
•When a covalent acid dissolves in water, the proton
transfer that forms H3O+ is called dissociation.
•When a strong acid dissolves in water, 100% of
the acid dissociates into ions.
HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
•A single reaction arrow can be used, because the
product is greatly favored at equilibrium.
•Common strong acids are HI, HBr, HCl, H2SO4,
and HNO3.
14
9.3 Acid and Base Strength
A. Relating Acid and Base Strength
•When a weak acid dissolves in water, only a
small fraction of the acid dissociates into ions.
•Unequal reaction arrows may be used, because the
reactants are usually favored at equilibrium.
CH3COOH(l) + H2O(l)
H3O+(aq) + CH3COO−(aq)
•Common weak acids are H3PO4, HF, H2CO3, and
HCN.
15
9.3 Acid and Base Strength
A. Relating Acid and Base Strength
A strong acid, HCl, is
completely dissociated
into H3O+(aq) and Cl−(aq).
A weak acid, CH3COOH,
contains mostly
undissociated acid.
16
9.3 Acid and Base Strength
A. Relating Acid and Base Strength
•A strong acid readily donates a proton, forming
a weak conjugate base.
HCl
strong acid
Cl−
weak conjugate base
•A strong base readily accepts a proton, forming
a weak conjugate acid.
OH−
strong base
H2O
weak conjugate acid
19
9.3 Acid and Base Strength
A. Relating Acid and Base Strength
9.2
-----
20
9.3 Acid and Base Strength
B. Using Acid Strength to Predict the Direction
of Equilibrium
•A Brønsted–Lowry acid–base reaction represents
an equilibrium.
H A
acid
+
B
base
A − +
H B+
conjugate conjugate
acid
base
•The position of the equilibrium depends upon the
strengths of the acids and bases.
•The stronger acid reacts with the stronger base
to form the weaker acid and the weaker base.
21
9.3 Acid and Base Strength
B. Using Acid Strength to Predict the Direction
of Equilibrium
•When the stronger acid and base are the reactants
on the left side, the reaction readily occurs and
the reaction proceeds to the right.
H A
+
B
stronger
stronger
acid
base
A − +
weaker
base
H B+
weaker
acid
•A larger forward arrow means that products are
favored.
22
9.3 Acid and Base Strength
B. Using Acid Strength to Predict the Direction
of Equilibrium
•If an acid–base reaction would form the stronger
acid and base, equilibrium favors the reactants
and little product forms.
H A
weaker
acid
+
B
weaker
base
A − +
stronger
base
H B+
stronger
acid
•A larger reverse arrow means that reactants are
favored.
23
9.3 Acid and Base Strength
HOW TO Predict the Direction of Equilibrium in an
Acid–Base Reaction
Are the reactants or products favored in
Example
the following acid–base reaction?
gain of H+
HCN(g)
acid
+
−OH(aq)
base
−CN(aq)
+ H2O(l)
conjugate conjugate
acid
base
loss of H+
Step [1]
Identify the acid in the reactants and the
conjugate acid in the products.
24
9.3 Acid and Base Strength
HOW TO Predict the Direction of Equilibrium in an
Acid–Base Reaction
Step [2]
Determine the relative strength of the acid
and the conjugate acid.
•From Table 9.2, HCN is a stronger acid than H2O.
Step [3]
HCN(g)
stronger
acid
Equilibrium favors the formation of the
weaker acid.
+
−OH(aq)
−CN(aq)
+ H2O(l)
weaker
acid
Products are favored.
25
9.4 Equilibrium and Acid Dissociation
Constants
For the reaction where an acid (HA) dissolves in
water,
HA(g) + H2O(l)
H3O+(aq) + A
− (aq)
the following equilibrium constant can be written:
K
=
[H3O+][ A −]
[HA][H2O]
26
9.4 Equilibrium and Acid Dissociation
Constants
•Multiplying both sides by [H2O] forms a new (more useful)
constant, called the acid dissociation constant, Ka.
Ka
=
K[H2O]
=
[H3O+][ A −]
[HA]
acid dissociation
constant
•The stronger the acid, the larger the Ka value.
•Equilibrium favors formation of the weaker acid,
the acid with the smaller Ka value.
27
9.4 Equilibrium and Acid Dissociation
Constants
28
9.5 Dissociation of Water
Water can behave as both a Brønsted–Lowry acid
and a Brønsted–Lowry base. Thus, two water
molecules can react together in an acid–base reaction:
loss of H+
H
O
acid
H + H
O
base
H
+
H
−
O
H
conjugate
base
+
H
O
H
conjugate
acid
gain of H+
29
9.5 Dissociation of Water
•From the reaction of two water molecules, the
following equilibrium constant expression can be
written:
K =
[H3O+][OH−]
[H2O]2
•Multiplying both sides by [H2O]2 yields Kw, the
ion-product constant for water.
Kw =
ion-product
constant
[H3O+][OH−]
30
9.5 Dissociation of Water
•Experimentally it can be shown that:
[H3O+] = [OH−] = 1.0 x 10−7 M at 25 oC
Kw
=
[H3O+] [OH−]
Kw
=
(1.0 x 10−7) x (1.0 x 10−7)
Kw
=
1.0 x 10−14
•Kw is a constant, 1.0 x 10−14, for all aqueous
solutions at 25 oC.
31
9.5 Dissociation of Water
If the [H3O+] in a cup of coffee is 1.0 x 10−5 M, then
the [OH−] can be calculated as follows:
[OH−] =
Kw
[H3O+]
1.0 x 10−14
=
=
−5
1.0 x 10
1.0 x 10−9 M
In this cup of coffee, therefore, [H3O+] > [OH–], and
the solution is acidic overall.
33
9.5 Dissociation of Water
34
9.6 The pH Scale
A. Calculating pH
The lower the pH, the higher the concentration of
H3O+:
•Acidic solution:
[H3O+] > 1 x 10−7
•Neutral solution:
[H3O+] = 1 x 10−7  pH = 7
•Basic solution:
[H3O+] < 1 x 10−7
 pH < 7
 pH > 7
35
9.6 The pH Scale
A. Calculating pH
36
9.6 The pH Scale
B. Calculating pH Using a Calculator
•A logarithm has the same number of places after
the decimal as there are digits in the original number.
•Example:
If [H3O+] = 1.2 x 10–5 M for a solution, what is its pH?
pH = –log [H3O+]
pH = –log (1.2 x 10–5) 2 digits
pH = –(–4.92) 2 decimal places
pH = 4.92
The solution is acidic because the pH < 7.
37
9.6 The pH Scale
B. Calculating pH Using a Calculator
If the pH of a solution is 8.50, what is the [H3O+]?
pH = −log [H3O+]
8.50 = −log [H3O+]
−8.50 = log [H3O+]
antilog (−8.50 ) = [H3O+]
2 decimal places
[H3O+] = 3.2 x 10−9 M
2 digits
The solution is basic because [H3O+] > 1 x 10–7 M.
38
9.6 The pH Scale
39
9.7 Common Acid–Base Reactions
A. Reaction of Acids with Hydroxide Bases
Neutralization reaction: An acid-base reaction that
produces a salt and water as products.
HA(aq) + MOH(aq)
base
acid
H
OH(l) + MA(aq)
water
salt
•The acid HA donates a proton (H+) to the OH− base
to form H2O.
•The anion A− from the acid combines with the
cation M+ from the base to form the salt MA.
40
9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction Between HA and MOH
Example
Write a balanced equation for the reaction
of Mg(OH)2 with HCl.
Step [1]
Identify the acid and base in the reactants
and draw H2O as one product.
HCl(aq) + Mg(OH)2(aq)
base
acid
H2O(l) + salt
water
41
9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction between HA and MOH
Step [2]
Determine the structure of the salt.
•The salt is formed from the parts of the acid and
base that are not used to form H2O.
HCl
Mg(OH)2
H+
Cl−
reacts to used to
form H2O form salt
Mg2+
2 OH−
used to
react to
form salt form water
Mg2+ and Cl− combine to form MgCl2.
42
9.7 Common Acid–Base Reactions
HOW TO Draw a Balanced Equation for a Neutralization
Reaction between HA and MOH
Step [3]
Balance the equation.
Place a 2 to
balance O and H.
2 HCl(aq) + Mg(OH)2(aq)
base
acid
2 H2O(l) + MgCl2
water
salt
Place a 2 to
balance Cl.
43
9.7 Common Acid–Base Reactions
A. Reaction of Acids with Hydroxide Bases
A net ionic equation contains only the species involved
in a reaction.
HCl(aq) + NaOH(aq)
H—OH(l) + NaCl(aq)
•Written as individual ions:
H+(aq) + Cl−(aq) + Na+(aq) + OH− (aq)
H—OH(l) + Na+(aq) + Cl−(aq)
•Omit the spectator ions, Na+ and Cl–.
•What remains is the net ionic equation:
H+(aq) + OH− (aq)
H—OH(l)
44
9.7 Common Acid–Base Reactions
B. Reaction of Acids with Bicarbonate Bases
•A bicarbonate base, HCO3−, reacts with one H+ to
form carbonic acid, H2CO3.
H+(aq) + HCO3−(aq)
H2CO3(aq)
H2O(l) + CO2(g)
•Carbonic acid then decomposes into H2O and CO2.
•For example:
HCl(aq) + NaHCO3(aq)
NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)
45
9.7 Common Acid–Base Reactions
B. Reaction of Acids with Bicarbonate Bases
•A carbonate base, CO32–, reacts with two H+ to
form carbonic acid, H2CO3.
2 H+(aq) + CO32–(aq)
H2CO3(aq)
H2O(l) + CO2(g)
•For example:
2 HCl(aq) + Na2CO3(aq)
2 NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)
46
9.8 The Acidity and Basicity of Salt
Solutions
A salt can form an acidic, basic, or neutral
solution depending on whether its cation and anion
are derived from a strong or weak acid and base.
For the salt M+A−:
•The cation M+ comes from the base.
•The anion A− comes from the acid HA.
47
9.8 The Acidity and Basicity of Salt
Solutions
NaCl
Na+
from NaOH
strong base
Cl−
from HCl
strong acid
A salt derived from a strong acid and strong base
forms a neutral solution (pH = 7).
48
9.8 The Acidity and Basicity of Salt
Solutions
NaHCO3
Na+
from NaOH
strong base
HCO3−
from H2CO3
weak acid
A salt derived from a strong base and a weak acid
forms a basic solution (pH > 7).
49
9.8 The Acidity and Basicity of Salt
Solutions
NH4Cl
NH4+
from NH3
weak base
Cl−
from HCl
strong acid
A salt derived from a weak base and a strong acid
forms an acidic solution (pH < 7).
50
9.8 The Acidity and Basicity of Salt
Solutions
Thus, the ion derived from the stronger acid or
base determines whether the solution is acidic
or basic.
51
9.9 Titration
•To determine the concentration of an acid or base
in a solution, we carry out a titration.
•If we want to know the concentration of an acid
solution, a base of known concentration is
added slowly until the acid is neutralized.
•When the acid is neutralized:
# of moles of acid = # of moles of base
•This is called the end point of the titration.
52
9.9 Titration
53
9.9 Titration
Determining an unknown molarity from titration data
requires three operations:
mole–mole
conversion
factor
Moles of
base
M (mol/L)
conversion
factor
[1]
Volume of
base
[2]
Moles of
acid
[3]
M (mol/L)
conversion
factor
Volume of
acid
54
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Example What is the molarity of an HCl solution if
22.5 mL of a 0.100 M NaOH solution are
needed to titrate a 25.0 mL sample of the
acid?
volume of base (NaOH)
22.5 mL
conc. of base (NaOH)
0.100 M
volume of acid (HCl)
25.0 mL
conc. of acid (HCl)
?
55
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Step [1]
Determine the number of moles of base
used to neutralize the acid.
Volume of
base
22.5 mL NaOH x
M (mol/L)
conversion
factor
1L
x 0.100 mol NaOH
1000 mL
1L
=
0.00225 mol NaOH
Moles of
base
56
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Step [2] Determine the number of moles of acid
that react from the balanced chemical
equation.
HCl(aq) + NaOH(aq)
H2O(l) + NaCl(aq)
mole–mole
conversion
Moles of
Moles of
factor
base
acid
0.00225 mol NaOH x 1 mol HCl = 0.00225 mol HCl
1 mol NaOH
57
9.9 Titration
HOW TO Determine the Molarity of an Acid Solution
from Titration
Step [3]
M
=
=
Determine the molarity of the acid from
the number of moles and the known
volume.
mol
L
=
0.00225 mol HCl x
25.0 mL solution
0.0900 M HCl
1000 mL
1L
mL—L
conversion
factor
3 sig. fig. Answer
58
9.10 Buffers
A buffer is a solution whose pH changes very little
when acid or base is added.
Most buffers are solutions composed of roughly
equal amounts of:
•A weak acid
•The salt of its conjugate base
The buffer resists change in pH because
•Added base, OH−, reacts with the weak acid
•Added acid, H3O+, reacts with the conjugate base
59
9.10 Buffers
A. General Characteristics of a Buffer
If an acid is added to the following buffer equilibrium,
then the excess acid reacts with the conjugate base,
so the overall pH does not change much.
60
9.10 Buffers
A. General Characteristics of a Buffer
If a base is added to the following buffer equilibrium,
then the excess base reacts with the conjugate acid,
so the overall pH does not change much.
61
9.10 Buffers
62
9.10 Buffers
B. Calculating the pH of a Buffer
•The effective pH range of a buffer depends on its Ka.
H3O+(aq) + A −(aq)
HA(aq) + H2O(l)
Ka =
[H3O+][ A −]
[HA]
•Rearranging this expression to solve for [H3O+]:
[H3
O+]
=
determines the
buffer pH
Ka
x
[HA]
[ A −]
63
9.11 Focus on the Human Body
Buffers in the Blood
•Normal blood pH is between 7.35 and 7.45.
•The principal buffer in the blood is carbonic acid/
bicarbonate (H2CO3/HCO3−).
CO2(g) + H2O(l)
H2CO3(aq)
H2O
H3O+(aq) + HCO3−(aq)
•CO2 is constantly produced by metabolic
processes in the body.
•The amount of CO2 is related to the pH of the blood.
64
9.11 Focus on the Human Body
Buffers in the Blood
Respiratory acidosis results when the body fails to
eliminate enough CO2, due to lung disease or failure.
65
9.11 Focus on the Human Body
Buffers in the Blood
Respiratory alkalosis is caused by hyperventilating;
very little CO2 is produced by the body.
66