Chapter 15

```Chapter 15
Acid–Base Equilibria
Chapter 15
15.1
15.2
15.3
15.4
15.5
Solutions of Acids or Bases Containing a
Common Ion
Buffered Solutions
Buffering Capacity
Titrations and pH Curves
Acid–Base Indicators
2
Section 15.1
Solutions of Acids or Bases Containing a Common Ion
Common Ion Effect
• Shift in equilibrium position that occurs because
equilibrium reaction.
• An application of Le Châtelier’s principle.
3
Section 15.1
Solutions of Acids or Bases Containing a Common Ion
Example
HCN(aq) + H2O(l)
H3O+(aq) + CN-(aq)
• Addition of NaCN will shift the equilibrium to the
left because of the addition of CN-, which is
already involved in the equilibrium reaction.
• A solution of HCN and NaCN is less acidic than
a solution of HCN alone.
4
Section 15.2
Atomic Masses
Buffered
Solutions
• Buffered Solution – resists a change in pH.
• They are weak acids or bases containing a
common ion.
• After addition of strong acid or base, deal with
stoichiometry first, then the equilibrium.
5
Section 15.2
Atomic Masses
Buffered
Solutions
Adding an Acid to a Buffer
6
Section 15.2
Atomic Masses
Buffered
Solutions
Buffers
7
Section 15.2
Atomic Masses
Buffered
Solutions
Solving Problems with Buffered Solutions
8
Section 15.2
Atomic Masses
Buffered
Solutions
Buffering: How Does It Work?
9
Section 15.2
Atomic Masses
Buffered
Solutions
Buffering: How Does It Work?
10
Section 15.2
Atomic Masses
Buffered
Solutions
Henderson–Hasselbalch Equation

pH = pK a
 A 
+ lo g
H A 
• For a particular buffering system (conjugate
acid–base pair), all solutions that have the
same ratio [A–] / [HA] will have the same pH.
11
Section 15.2
Atomic Masses
Buffered
Solutions
Exercise
What is the pH of a buffer solution that is 0.45
M acetic acid (HC2H3O2) and 0.85 M sodium
acetate (NaC2H3O2)? The Ka for acetic acid is
1.8 × 10–5.
pH = 5.02
12
Section 15.2
Atomic Masses
Buffered
Solutions
13
Section 15.2
Atomic Masses
Buffered
Solutions
Buffered Solution Characteristics
• Buffers contain relatively large amounts of weak
acid and corresponding conjugate base.
• Added H+ reacts to completion with the
conjugate base.
• Added OH reacts to completion with the weak
acid.
14
Section 15.2
Atomic Masses
Buffered
Solutions
Buffered Solution Characteristics
• The pH in the buffered solution is determined by
the ratio of the concentrations of the weak acid
and weak base. As long as this ratio remains
virtually constant, the pH will remain virtually
constant. This will be the case as long as the
concentrations of the buffering materials (HA
and A– or B and BH+) are large compared with
amounts of H+ or OH– added.
15
Section 15.3
The Mole Capacity
Buffering
• The amount of protons or hydroxide ions the
buffer can absorb without a significant change
in pH.
• Determined by the magnitudes of [HA] and [A–].
• A buffer with large capacity contains large
concentrations of the buffering components.
16
Section 15.3
The Mole Capacity
Buffering
• Optimal buffering occurs when [HA] is equal to
[A–].
• It is for this condition that the ratio [A–] / [HA] is
most resistant to change when H+ or OH– is
17
Section 15.3
The Mole Capacity
Buffering
Choosing a Buffer
• pKa of the weak acid to be used in the buffer
should be as close as possible to the desired
pH.
18
Section 15.4
Titrations and pH Curves
Titration Curve
• Plotting the pH of the solution being analyzed
as a function of the amount of titrant added.
• Equivalence (Stoichiometric) Point – point in the
titration when enough titrant has been added to
react exactly with the substance in solution
being titrated.
19
Section 15.4
Titrations and pH Curves
Neutralization of a Strong Acid with a Strong Base
20
Section 15.4
Titrations and pH Curves
The pH Curve for the Titration of 50.0 mL of 0.200 M HNO3 with
0.100 M NaOH
21
Section 15.4
Titrations and pH Curves
The pH Curve for the Titration of 100.0 mL of 0.50 M NaOH
with 1.0 M HCI
22
Section 15.4
Titrations and pH Curves
Weak Acid–Strong Base Titration
Step 1:
Step 2:
A stoichiometry problem (reaction is
assumed to run to completion) then
determine remaining species.
An equilibrium problem (determine
position of weak acid equilibrium and
calculate pH).
23
Section 15.4
Titrations and pH Curves
Concept Check
Consider a solution made by mixing 0.10 mol of
HCN (Ka = 6.2 x 10–10) with 0.040 mol NaOH in
1.0 L of aqueous solution.
What are the major species immediately upon
mixing (that is, before a reaction)?
HCN, Na+, OH–, H2O
24
Section 15.4
Titrations and pH Curves
• Why isn’t NaOH a major species?
• Why aren’t H+ and CN– major species?
• List all possibilities for the dominant reaction.
25
Section 15.4
Titrations and pH Curves
The possibilities for the dominant reaction are:
1.
2.
3.
4.
5.
H2O(l) + H2O(l)
H3O+(aq) + OH–(aq)
HCN(aq) + H2O(l)
H3O+(aq) + CN–(aq)
HCN(aq) + OH–(aq)
CN–(aq) + H2O(l)
Na+(aq) + OH–(aq)
NaOH
Na+(aq) + H2O(l)
NaOH + H+(aq)
26
Section 15.4
Titrations and pH Curves
• How do we decide which reaction controls the
pH?
H2O(l) + H2O(l)
H3O+(aq) + OH–(aq)
HCN(aq) + H2O(l)
H3O+(aq) + CN–(aq)
HCN(aq) + OH–(aq)
CN–(aq) + H2O(l)
27
Section 15.4
Titrations and pH Curves
HCN(aq) + OH–(aq)
CN–(aq) + H2O(l)
• What are the major species after this reaction
occurs?
HCN, CN–, H2O, Na+
28
Section 15.4
Titrations and pH Curves
• Now you can treat this situation as before.
• List the possibilities for the dominant reaction.
• Determine which controls the pH.
29
Section 15.4
Titrations and pH Curves
Concept Check
Calculate the pH of a solution made by mixing
0.20 mol HC2H3O2 (Ka = 1.8 x 10–5) with 0.030
mol NaOH in 1.0 L of aqueous solution.
30
Section 15.4
Titrations and pH Curves
• What are the major species in solution?
Na+, OH–, HC2H3O2, H2O
• Why isn’t NaOH a major species?
• Why aren’t H+ and C2H3O2– major species?
31
Section 15.4
Titrations and pH Curves
• What are the possibilities for the dominant
reaction?
1.
2.
3.
4.
5.
H2O(l) + H2O(l)
H3O+(aq) + OH–(aq)
HC2H3O2(aq) + H2O(l)
H3O+(aq) + C2H3O2–(aq)
HC2H3O2(aq) + OH–(aq)
C2H3O2–(aq) + H2O(l)
Na+(aq) + OH–(aq)
NaOH
Na+(aq) + H2O(l)
NaOH + H+(aq)
• Which of these reactions really occur?
32
Section 15.4
Titrations and pH Curves
• Which reaction controls the pH?
H2O(l) + H2O(l)
H3O+(aq) + OH–(aq)
HC2H3O2(aq) + H2O(l)
H3O+(aq) + C2H3O2–(aq)
HC2H3O2(aq) + OH–(aq)
C2H3O2–(aq) + H2O(l)
• How do you know?
33
Section 15.4
Titrations and pH Curves
HC2H3O2(aq) + OH– 
Before
Change
After
0.20 mol
0.030 mol
–0.030 mol –0.030 mol
0.17 mol
0
C2H3O2–(aq) + H2O
0
+0.030 mol
0.030 mol
K = 1.8 x 109
34
Section 15.4
Titrations and pH Curves
Steps Toward Solving for pH
H3O+
+ C2H3O2-(aq)
0.170 M
~0
0.030 M
–x
+x
+x
0.170 – x
x
0.030 + x
HC2H3O2(aq) + H2O
Initial
Change
Equilibrium
Ka = 1.8 x 10–5
pH = 3.99
35
Section 15.4
Titrations and pH Curves
Exercise
Calculate the pH of a 100.0 mL solution of
0.100 M acetic acid (HC2H3O2), which has a Ka
value of 1.8 x 10–5.
pH = 2.87
36
Section 15.4
Titrations and pH Curves
Concept Check
Calculate the pH of a solution made by mixing
100.0 mL of a 0.100 M solution of acetic acid
(HC2H3O2), which has a Ka value of 1.8 x 10–5,
and 50.0 mL of a 0.10 M NaOH solution.
pH = 4.74
37
Section 15.4
Titrations and pH Curves
Concept Check
Calculate the pH of a solution at the equivalence
point when 100.0 mL of a 0.100 M solution of
acetic acid (HC2H3O2), which has a Ka value of
1.8 x 10–5, is titrated with a 0.10 M NaOH
solution.
pH = 8.72
38
Section 15.4
Titrations and pH Curves
The pH Curve for the Titration of 50.0 mL of 0.100 M HC2H3O2 with
0.100 M NaOH
39
Section 15.4
Titrations and pH Curves
The pH Curves for
the Titrations of
50.0-mL Samples of
0.10 M Acids with
Various Ka Values
with 0.10 M NaOH
40
Section 15.4
Titrations and pH Curves
The pH Curve for the Titration of 100.0mL of 0.050 M NH3 with
0.10 M HCl
41
Section 15.5
Acid–Base Indicators
• Marks the end point of a titration by changing
color.
• The equivalence point is not necessarily the
same as the end point (but they are ideally as
close as possible).
42
Section 15.5
Acid–Base Indicators
The Acid and Base
Forms of the
Indicator
Phenolphthalein
43
Section 15.5
Acid–Base Indicators
The Methyl Orange Indicator is Yellow in Basic Solution and Red
in Acidic Solution
44
Section 15.5
Acid–Base Indicators
Useful pH Ranges for Several Common Indicators
45
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