Chapter 17
A Review of Strong Electrolytes
Ionic Equilibria: Acids and Bases
We must distinguish weak acids and bases
from strong electrolytes.
Weak acids and bases ionize or dissociate
much less than 100%.
In this chapter we will see that it is often less
than 10%!
Strong electrolytes ionize or dissociate
completely.
Strong electrolytes approach 100% dissociation
in aqueous solutions.
1
A Review of Strong Electrolytes
1
2
A Review of Strong Electrolytes
There are three classes of strong electrolytes.
Strong Water Soluble Acids
2
Strong Water Soluble Bases
The entire list of these bases was also introduced in
Chapter 4.
Remember the list of strong acids from Chapter 4.
+
2 O ≈100%
KOH (s) H
→ K (aq)
+ OH -(aq)
+
−
100%
HNO3(l) + H2O(l) ≈
→ H3O(aq)
+ NO3(aq)
2+
2 O ≈100%
Sr(OH) 2(s) H
→ Sr(aq)
+ 2 OH -(aq)
or
≈100%
+
−
HNO3(l) → H(aq)
+ NO3(aq)
3
4
A Review of Strong Electrolytes
3
ACID-BASE THEORIES
ACID
ACID-BASE
Most Water Soluble Salts
The solubility guidelines from Chapter 4 will help you remember
these salts.
+
2 O ≈100%
NaCl (s) H
→ Na (aq)
+ Cl -(aq)
The most general theory for common
aqueous acids and bases is the
BRØ
BRØNSTED - LOWRY theory
ACIDS
2+
−
2 O ≈100%
Ca(NO 3 ) 2(s ) H
→ Ca (aq)
+ 2 NO3(aq)
DONATE H+ IONS
BASES
5
ACCEPT H+ IONS
6
1
ACID-BASE THEORIES
ACID
ACID-BASE
ACID-BASE THEORIES
ACID
ACID-BASE
NH3 is a BASE in water — and water is
itself an ACID
The Brø
Brønsted definition means NH3 is a
BASE in water — and water is itself an
ACID
NH3
Base
+ H 2O
Acid
NH4+ + OH Acid
Base
NH3 / NH4+ is a conjugate pair — related by the gain
or loss of H+
Every acid has a conjugate base - and vicevice-versa.
7
8
More About Water
Conjugate Pairs
H2O can function as both an ACID and a BASE.
In pure water there can be AUTOIONIZATION
Equilibrium constant for autoion = Kw
Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC
9
The Autoionization of Water
HO
The pH and pOH scales
We can write the autoionization of water as a
dissociation reaction similar to those previously done
in this chapter.
+
2 (l)
2 (l )
3 (aq)
(aq)
+H O
→H O
←
+ OH
Because the activity of pure water is 1, the
equilibrium constant for this reaction is:
[
][
10
K c = H 3O + OH −
]
11
A convenient way to express the acidity and basicity
of a solution is the pH and pOH scales.
The pH of an aqueous solution is defined as:
[
pH = -log H 3O +
]
12
2
The pH and pOH scales
The pH and pOH scales
A convenient relationship between pH and pOH may
be derived for all dilute aqueous solutions at 250C.
+
−
[H3O ][OH ] = 1.0 ×10
−14
[
]
[
] (
[
])
- log H 3 O + + − log OH − = 14 .00
Taking the logarithm of both sides of this equation
gives:
[
Multiplying both sides of this equation by -1 gives:
Which can be rearranged to this form:
pH + pOH = 14.00
]
log H 3 O + + log OH − = − 14 . 00
13
The pH and pOH scales
14
The Autoionization of Water
Remember these two expressions!!
Example: Calculate the concentrations of H3O+
and OH- in 0.050 M HCl.
[H O ][OH ] =1.0×10
+
−
−14
3
pH+ pOH=14.00
15
Calculating [H3O+] & [OH-]
16
[H3O+], [OH-] and pH
You add 0.0010 mol of NaOH to 1.0 L of
pure water. Calculate [H3O+] and [OH-].
What is the pH of the
0.0010 M NaOH solution?
General conclusion —
Basic solution pH > 7
Neutral
pH = 7
Acidic solution
pH < 7
17
18
3
[H33O++], [OH--] and pH
The pH Scale
If the pH of Coke is 3.12, it is
____________.
Active Figure 17.2
19
Ionization Constants for Weak Monoprotic Acids and Bases
The pH and pOH scales
To help develop familiarity with the pH and pOH scale we can
look at a series of solutions in which [H3O+] varies between 1.0
M and 1.0 x 10-14 M.
[H3O+]
1.0 M
[OH-]
1.0 x 10-14 M
pH
0.00
pOH
14.00
1.0 x 10-3 M
1.0 x 10-11 M
3.00
11.00
10-7 M
1.0 x 10-7 M
7.00
7.00
2.0 x 10-12 M
5.0 x 10-3 M
11.70
2.30
1.0 M
14.00
0.00
1.0 x
1.0 x
10-14 M
20
Let’s look at the dissolution of acetic acid, a weak acid,
in water as an example.
The equation for the ionization of acetic acid is:
The equilibrium constant for this ionization is
expressed as:
+
Kc =
[H O ][CH COO ]
3
−
3
[CH 3COOH ]
21
Ionization Constants for Weak
Monoprotic Acids and Bases
Ionization Constants for Weak
Monoprotic Acids and Bases
We can define a new equilibrium constant for weak
acid equilibria that uses the previous definition.
This equilibrium constant is called the acid ionization
constant.
The symbol for the ionization constant is Ka.
[H O ][CH COO ] = 1.8 ×10
+
Ka =
3
The ionization constant values for several acids are
given below.
−
3
22
−5
[CH 3COOH ]
for acetic acid
23
Which acid is the strongest?
Acid
Formula
Ka value
Acetic
CH3COOH
1.8 x 10-5
Nitrous
HNO2
4.5 x 10-4
Hydrofluoric
HF
7.2 x 10-4
Hypochlorous
HClO
3.5 x 10-8
Hydrocyanic
HCN
4.0 x 10-10
24
4
Ionization Constants for Weak
Monoprotic Acids and Bases
Ionization Constants for Weak
Monoprotic Acids and Bases
From the above table we see that the order of
increasing acid strength for these weak acids is:
HF > HNO2 > CH3COOH > HClO > HCN
Example: Write the equation for the ionization of the
weak acid HCN and the expression for its ionization
constant.
→ H + + CN HCN ←
The order of increasing base strength of the anions
(conjugate bases) of these acids is:
[H ][CN ] = 4.0 x 10
+
Ka =
F- < NO-2 < CH3COO- < ClO- < CN-
-
-10
[HCN ]
25
Ionization Constants for Weak
Monoprotic Acids and Bases
26
Ionization Constants for Weak
Monoprotic Acids and Bases
Example: The pH of a 0.10 M solution of a weak
monoprotic acid, HA, is found to be 2.97. What is
the value for its ionization constant?
Example: Calculate the concentrations of the
various species in 0.15 M acetic acid, CH3COOH,
solution.
It is always a good idea to write down the ionization
reaction and the ionization constant expression.
+ H 2O →
← H 3 O + + CH 3 COO
CH 3 COOH
K
a
=
[H
3
O
+
[CH
][CH
3
3 COO
COOH ]
-
] = 1 . 8 × 10
-
−5
Do problem
27
Ionization Constants for Weak
Monoprotic Acids and Bases
Equilibria Involving A Weak Acid
Consider the approximate expression
K a = 1.8 x 10 -5 =
x2
1.00
28
x = [H3 O+ ] = [K a • 1.00]1/2
For many weak acids
Example: Calculate the concentrations of the
species in 0.15 M hydrocyanic acid, HCN,
solution.
Ka= 4.0 x 10-10 for HCN
[H3O+] = [conj. base] = [Ka • Co]1/2
where C0 = initial conc. of acid
Useful Rule of Thumb:
If 100•
100•Ka < Co, then [H3O+] = [Ka•Co]1/2
29
30
5
Equilibria
Equilibria Involving
Involving A
A Weak
Weak Acid
Acid
Ionization Constants for Weak
Monoprotic Acids and Bases
Calculate the pH of a 0.0010 M solution of formic acid,
HCO2H.
HCO2H + H2O
HCO2- + H3O+
4
Ka = 1.8 x 10
Let’s look at the percent ionization of two weak
acids as a function of their ionization constants.
Solution
Ka
[H+]
pH
% ionization
0.15 M
acetic acid
1.8 x 10-5
1.6 x 10-3
2.80
1.1
0.15 M
HCN
4.0 x 10-10
7.7 x 10-6
5.11
0.0051
Note that the [H+] in 0.15 M acetic acid is 210
times greater than for 0.15 M HCN.
31
Ionization Constants for Weak
Monoprotic Acids and Bases
All of the calculations and understanding we have at present
can be applied to weak acids and weak bases!
Example: Calculate the concentrations of the various species
in 0.10 M aqueous ammonia and the pH.
32
Equilibria Involving A Weak Base
NH3 + H2O
Kb = 1.8 x 10-5
NH4+ +
OH-
33
Equilibrium Constants
for Weak Acids
34
Equilibrium Constants
for Weak Bases
Weak acid has Ka < 1
Leads to small [H3O+] and a pH of 2 - 7
Weak base has Kb < 1
Leads to small [OH-] and a pH of 12 - 7
35
36
6
Ionization Constants for Acids/Bases
Conjugate
Bases
Acids
Relation
Increase
strength
of Ka, Kb,
[H3O+]
and pH
Increase
strength
37
K and Acid-Base Reactions
Acid
Acid-Base
K and AcidAcid-Base Reactions
ACIDS
ACIDS
STRONG
STRONG
A strong acid is 100% dissociated.
CONJUGATE
CONJUGATE BASES
BASES
weak
weak
weak
weak
38
Therefore, a STRONG ACID—
ACID—a good H+
donor—
donor—must have a WEAK CONJUGATE
BASE—
BASE—a poor H+ acceptor.
HNO3(aq) + H2O(liq) H3O+(aq) + NO3-(aq)
STRONG
STRONG
Reactions always go from the stronger
A-B pair (larger K) to the weaker AA-B
pair (smaller K).
STRONG A
base
acid
weak B
••Every
Every A-B reaction
A
A-B
reaction has
has two
two acids
acids and
and two
two bases.
bases.
••Equilibrium
Equilibrium always
always lies
lies toward
toward the
the weaker
weaker pair.
pair.
••Here
Here K
K is
is very
very large.
large.
39
K and Acid-Base Reactions
Acid
Acid-Base
BASE
HNO3 + H2O
STRONG
ACID
40
K and Acid-Base Reactions
Acid
Acid-Base
Acetic acid is only 0.42% ionized when [HOAc
[HOAc]] = 1.0 M. It is
a WEAK ACID
ACID
H3O+ + NO3WEAK
BASE
HOAc
WEAK A
+
H3O+
H2O
base
acid
+
OAc-
STRONG B
Because [H3O+] is small, this must mean
We know from experiment that HNO3
is a strong acid.
1. It is a stronger acid than H3O+
2. H2O is a stronger base than NO33. K for this reaction is large
1. H3O+ is a stronger acid than HOAc
2. OAc- is a stronger base than H2O
3. K for this reaction is small
41
42
7
Types of Acid/Base Reactions
Types of Acid/Base Reactions
Weak acid (acetic ac.) + Strong base (NaOH
(NaOH))
Strong acid (HCl
(HCl)) + Strong base (NaOH
(NaOH))
H+ + Cl- + Na+ + OH-
H2O + Na+ + ClCH3CO2H + OHH2O + CH3CO2 This is the reverse of the reaction of CH3CO2(conjugate base) with H2O.
OH- stronger base than CH3CO2 K = 1/Kb = 1/(5.6 x 10-10) = 1.8 x 109
Net ionic equation
H+(aq)
(aq)
+
OH-(aq)
aq)
H2O(liq)
K = 1/Kw = 1 x 1014
Mixing equal molar quantities of a strong
acid and strong base produces a neutral
solution.
Mixing
Mixing equal
equal molar
molar quantities
quantities of
of aa
weak
weak acid
acid and
and strong
strong base
base produces
produces
the
the acid’s
acid’s conjugate
conjugate base.
base. The
The
solution
solution is
is basic.
basic.
43
44
Polyprotic Acids
Types of Acid/Base Reactions
Strong acid (HCl
(HCl)) + Weak base (NH3)
H3O+ + NH3
H2O + NH4+
This is the reverse of the reaction of NH4+
(conjugate acid of NH3) with H2O.
H3O+ stronger acid than NH4+
K = 1/Ka = 1.8 x 109
Many weak acids contain two or more acidic hydrogens.
The calculation of equilibria for polyprotic acids is done
in a stepwise fashion.
Consider arsenic acid, H3AsO4, which has three
ionization constants.
Ka1 = 2.5 x 10-4
Ka2 = 5.6 x 10-8
Ka3 = 3.0 x 10-13
Mixing
Mixing equal
equal molar
molar quantities
quantities of
of aa
strong
strong acid
acid and
and weak
weak base
base produces
produces
the
the bases’s
bases’s conjugate
conjugate acid.
acid. The
The
solution
solution is
is acid.
acid.
1
2
3
Examples include H3PO4 and H3AsO4.
There is an ionization constant for each step.
45
Polyprotic Acids
46
Polyprotic Acids
1
Example: Calculate the concentration of all species
in 0.100 M arsenic acid, H3AsO4, solution.
Write the first ionization step and represent the
concentrations.
Species Concentration
Approach this problem exactly as previously done.
H3AsO4 →
← H + + H 2 AsO−4
(0.100 − x )M
xM
A comparison of the various species in 0.100 M
H3AsO4 solution follows.
xM
47
H3AsO4
0.095 M
H+
0.0049 M
H2AsO4-
0.0049 M
HAsO42-
5.6 x 10-8 M
AsO43-
3.4 x 10-18 M
OH-
2.0 x 10-12 M
48
8
AcidAcid-Base Properties of Salts
AcidAcid-Base Properties of Salts
Calculate the pH of a 0.10 M solution of
Na2CO3.
Na+ + H2O --->
---> neutral
CO32- + H2O
HCO3- + OHbase
acid
acid
base
Kb = 2.1 x 10-4
49
Salts of Weak Bases and Strong Acids
50
Salts of Weak Bases and Weak Acids
Example: Calculate [H+], pH, and percent
hydrolysis for the ammonium ion in 0.10 M
ammonium bromide, NH4Br, solution.
The fluoride ion hydrolyzes to produce OH- ions. Its
hydrolysis constant is (base dissociation):
F− + H 2O →
← HF + OH −
Kb =
[ HF][OH − ]
[ ]
F-
=
Kw
Ka for HF
10
. × 10−14
Kb =
= 14
. × 10−11
7.2 × 10−4
Because the Ka for (CH3)3NH+ ions is one order of
magnitude larger than the Kb for F- ions, H+ ions are
produced in excess making the solution acidic.
51
Salts of Weak Bases and Weak Acids
1
Salts of Weak Bases and Weak Acids
Summary of the major points of
hydrolysis up to now.
The reactions of anions of weak monoprotic
acids (from a salt) with water to form free
molecular acids and OH-.
A - + H 2O →
← HA + OH -
Kb =
52
2.
The reactions of anions of weak monoprotic acids
(from a salt) with water to form free molecular acids
and OH-.
→ B + H O+
BH + + H 2O ←
3
K
K a = w B = weak base
K b ( B)
Kw
Ka ( HA )
53
54
9
Lewis Acids & Bases
Salts of Weak Bases and Weak Acids
Aqueous solutions of salts of strong acids
and strong bases are neutral.
Aqueous solutions of salts of strong bases
and weak acids are basic.
Aqueous solutions of salts of weak bases and
strong acids are acidic.
Aqueous solutions of salts of weak bases and
weak acids can be neutral, basic or acidic.
The values of Ka and Kb determine the pH.
Lewis acid
a substance that
accepts an electron
pair
Lewis base
a substance that
donates an electron
pair
55
Lewis Acids & Bases
Reaction of a Lewis Acid and
Lewis Base
56
Formation of hydronium ion is also an
excellent example.
New bond formed
using electron pair
from the Lewis
base.
Coordinate
covalent bond
H
+
ACID
••
•• O—H
••
H O—H
H
BASE
H
•Electron pair of the new OO-H bond
originates on the Lewis base.
57
58
Lewis Acids & Bases
Lewis Acid/Base Reaction
Other good examples involve metal
ions.
•• ••
O—H
••
Co2+
ACID
59
Co2+
H
BASE
•• O—H
H
60
10
Lewis Acids & Bases
Lewis AcidAcid-Base
Interactions in Biology
The combination of metal ions (Lewis acids)
with Lewis bases such as H2O and NH3
------>
------> COMPLEX IONS
The heme group in
hemoglobin can interact
with O2 and CO.
The Fe ion in hemoglobin
is a Lewis acid
O2 and CO can act as
Lewis bases
Heme group
61
Why?
62
Why is CH3CO2H an Acid?
Why are some compounds
acids?
Why are some compounds
bases?
Why do acids and bases
vary in strength?
Can we predict variations in
acidity or basicity?
basicity?
1. The electronegativity of the O atoms causes the H
attached to O to be highly positive.
2. The O—
O—H bond is highly polar.
3. The H atom of O—
O—H is readily attracted to polar H2O.
Figure 17.9
63
64
Basicity of Oxoanions
NO3Acetic acid
Ka = 1.8 x 10-5
CO32-
PO43-
Trichloroacetic acid
Ka = 0.3
Trichloroacetic acid is a much stronger acid owing to
the high electronegativity of Cl.
Cl withdraws electrons from the rest of the molecule.
This makes the O—
O—H bond highly polar. The H of O—
O—
H is very positive.
These ions are BASES.
They become more and more basic as the negative
charge increases.
As the charge goes up, they interact more strongly
with polar water molecules.
65
66
11