Unit 8
Ionic Equilibria I:
Acids and Bases
1
Chapter Goals
1.
2.
3.
4.
Identify strong electrolytes and calculate
concentrations of their ions
Understand the autoionization of water
Understand the pH and pOH scales
Perform calculations involving pH, pOH, kw.
2
A Review of Strong Electrolytes
RECALL:• Weak acids and bases ionize or dissociate
partially, 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.
3
A Review of Strong Electrolytes
• There are three classes of strong
electrolytes.
1 Strong Water Soluble Acids
Remember the list of strong acids from Chapter 4.
100%
HNO3( )  H 2O (  )  H 3O

(aq)

3(aq)
 NO
or
100%
HNO3( )  H

(aq)

3(aq)
 NO
4
A Review of Strong Electrolytes
100%
HNO3( )  H 2O (  )  H 3O

(aq)

3(aq)
 NO
or
100%
HNO3( )  H

(aq)

3(aq)
 NO
5
A Review of Strong Electrolytes
2 Strong Water Soluble Bases
The entire list of these bases was also introduced in
Chapter 4.
H 2 O 100%
KOH(s)   K
H 2 O 100%

(aq)
 OH
2
(aq)
Sr(OH) 2(s)   Sr
(aq)
 2 OH
(aq)
6
A Review of Strong Electrolytes
3 Most Water Soluble Salts
The solubility guidelines from Chapter 4 will help you
remember these salts.
H 2 O 100%
NaCl(s)   Na
H 2 O 100%

(aq)
 Cl
Ca(NO3 ) 2s    Ca
2
(aq)
(aq)

3(aq)
 2 NO
7
A Review of Strong Electrolytes
• The calculation of ion concentrations in solutions
of strong electrolytes is easy.
Example 18-1: Calculate the concentrations of ions
in 0.050 M nitric acid, HNO3.

(aq)
HNO3( )  H2O( ) 
 H3O
100%
0.050 M

3(aq)
 NO
0.050 M 0.050 M
8
A Review of Strong Electrolytes
Example 18-2: Calculate the concentrations of ions
in 0.020 M strontium hydroxide, Sr(OH)2, solution.
You do it!
2
(aq)
Sr(OH)2(s) 
Sr
H2 O
0.020M
 2 OH
(aq)
0.020M 20.020M 
0.040M
9
The Autoionization of Water
• Pure water ionizes very slightly.
– The concentration of the ionized water is less than
one-millionth molar at room temperature.
10
The Autoionization of Water
• We can write the autoionization of water as a
dissociation reaction similar to those previously
done in this chapter.


H 2O( )  H 2O()  H3O(aq)  OH(aq)
• Because the activity of pure water is 1, the
equilibrium constant for this reaction is:

Kc  H 3O
+

OH


11
The Autoionization of Water
• Experimental measurements have determined
that the concentration of each ion is 1.0 x 10-7 M
at 25oC.
– Note that this is at 25oC, not every temperature!
• We can determine the value of Kc from this
information.

OH 
 1.0 x 10 1.0 x 10 
Kc  H 3O

+
-7
 1.0 x10
-7
14
12
The Autoionization of Water
• This particular equilibrium constant is called the
ion-product for water and given the symbol Kw.
– Kw is one of the recurring expressions for the
remainder of this chapter and Chapters 19 and 20.

K w  H 3O
+
OH 
 1.0 x10

14
13
The Autoionization of Water
Example 18-3: Calculate the concentrations of H3O+
and OH- in 0.050 M HCl.
HCl + H 2O  H3O+  Cl
0.050M


Thus the
0.050M 0.050M
H O   0.050M .
+
3
The H 3O + and K w will allow us to calculate [OH - ].
14
The pH and pOH scales
• 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
+

15
The pH and pOH scales
• In general, a lower case p before a symbol is
read as the ‘negative logarithm of’ the symbol.
• Thus we can write the following notations.
 
pAg = -logAg 
pOH = -log OH
-
+
and so forth for other quantities .
16
The pH and pOH scales
• If either the [H3O+] or [OH-] is known, the pH and
pOH can be calculated.
Example 18-4: Calculate the pH of a solution in
which the [H3O+] =0.030 M.
pH = -log H O 
pH   log 3.0 10 
+
3
2
pH  1.52
17
The pH and pOH scales
Example 18-5: The pH of a solution is 4.597. What
is the concentration of H3O+?
You do it!

pH  -log[H 3O ]
4.597  -log[H 3O  ]
log[H 3O  ]  -4.597
[H 3O  ]  10 -4.597
[H 3O  ]  2.53  10 5 M
18
The pH and pOH scales
• A convenient relationship between pH and pOH
may be derived for all dilute aqueous solutions at
250C.


[H 3O ][OH ]  1.0 10
14
• Taking the logarithm of both sides of this
equation gives:






log H 3O  log OH  14.00
19
The pH and pOH scales
• Multiplying both sides of this equation by -1
gives:



- log H 3O   log OH


  14.00
• Which can be rearranged to this form:
pH  pOH  14.00
20
The pH and pOH scales
• Remember these two expressions!!
– They are key to the next three chapters!
H O OH   1.0 10


14
3
pH  pOH  14.00
21
The pH and pOH scales
• The usual range for the pH scale is 0 to 14.
H O   1.0 M to H O   1.0 10


3
14
3
pH  0
M
pH  14.00
to
• And for pOH the scale is also 0 to 14 but inverted
from pH.
– pH = 0 has a pOH = 14 and pH = 14 has a pOH = 0.
OH   1.0 10

pOH  14.00
14



M up to OH  1.0M
pOH  0
22
The pH and pOH scales
Example 18-6: Calculate the [H3O+], pH, [OH-], and
pOH for a 0.020 M HNO3 solution.
– Is HNO3 a weak or strong acid?
– What is the [H3O+] ?
100%
HNO 3  H 2 O 
 H 3O   NO3-
0.020M
0.020 M 0.020 M
H O   2.0 10 M
pH  -log 2.0 10 M 

2
3
2
pH  1.70
23
The pH and pOH scales
Example 18-6: Calculate the [H3O+], pH, [OH-], and
pOH for a 0.020 M HNO3 solution.
  
1.0 10
1.0 10
OH   H O   2.0 10  5.0 10
pOH   log 5.0 10   12.30
K w  H 3O  OH   1.0 10 14
14

14

2
13
M
3
13
24
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
1.0 x 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 x 10-14 M
1.0 M
14.00
0.00
25