Download pH/pOH

pH and pOH
Ionization of water

Experiments have shown that pure
water ionizes very slightly:



2H2O  H3O+ + OH-
Measurements show that:
[H3O+] = [OH-]=1 x 10-7 M
Pure water contains equal
concentrations of H3O+ + OH-, so it is
neutral.
pH

pH is a measure of the
concentration of hydronium
ions in a solution.

pH = -log [H3O+]
or

pH = -log [H+]
Sig. Figs. for Logarithms

The rule is that the number of decimal
places in the log is equal to the number
of significant figures in the original
number.

Example:


[H+] = 1.0 x 10-9 M (2 significant figures)
pH = -log(1.0 x 10-9) = 9.00 (2 decimal places)
Example: What is the pH of a solution
where [H3O+] = 1 x 10-7 M?
pH = -log [H3O+]
-7
 pH = -log(1 x 10 )
 pH = 7.0

Example: What is the pH of a solution
where [H3O+] = 1 x 10-5 M?
pH = -log [H3O+]
-5
 pH = -log(1 x 10 )
 pH = 5.0


When acid is added to water, the [H3O+]
increases, and the pH decreases.
Example: What is the pH of a solution
where [H3O+] = 1 x 10-10 M?
pH = -log [H3O+]
-10)
 pH = -log(1 x 10
 pH = 10.0


When base is added to water, the [H3O+]
decreases, and the pH increases.
The pH Scale
0
Acid
7
Neutral
14
Base
pOH

pOH is a measure of the
concentration of hydroxide
ions in a solution.

pOH = -log [OH-]
Example: What is the pOH of a solution
where [OH-] = 1 x 10-5 M?
pOH = -log [OH-]
-5
 pOH = -log(1 x 10 )
 pOH = 5.0

How are pH and pOH related?

At every pH, the following relationships
hold true:

[H+] • [OH-] = 1 x 10-14 M

pH + pOH = 14
Example 1: What is the pH of a solution
where [H+] = 3.4 x 10-5 M?
pH = -log [H+]
-5
 pH = -log(3.4 x 10
M)
 pH = 4.47

Example 2: The pH of a solution is
measured to be 8.86. What is the [H+] in
this solution?
pH = -log [H+]
+
 8.86 = -log [H ]
+
 -8.86 = log [H ]
+
 [H ] = antilog (-8.86)
+
-8.86
 [H ] = 10
+
-9 M
 [H ] = 1.4 x 10

Example 3: What is the pH of a solution
where [H+] = 5.4 x 10-6 M?
pH = -log [H+]
-6
 pH = -log(5.4 x 10 )
 pH = 5.27

Example 4: What is the [OH-] and pOH
for the solution in example #3?
[H3O+][OH-]= 1 x 10-14
-6
-14
 (5.4 x 10 )[OH ] = 1 x 10
-9 M
 [OH ] = 1.9 x 10

pH + pOH = 14
 pOH = 14 – 5.27 = 8.73

Buffered Solutions
A solution of a weak acid
and a common ion is called a
buffered solution.
Consider the following
buffered solution…
HAc  H+ + AcH2O  H+ + OHAdd additional acid…(H+)
The H+ will combine with the Ac- producing
HAc. There is an excess of Ac- from the common
ion salt.
HAc  H+ + Ac-
Now, add additional base (OH-)
The OH- will combine with the
H+ to produce water…
H2O  H+ + OHThe H+ comes from the HAc
HAc  H+ + Ac-
Thus, the solution maintains
it’s pH in spite of added acid or
base.