Acid-Base
Titration &
pH
16-1 Objectives
1. Describe the self-ionization of water
2. Define pH and give the pH of a neutral
solution at 25oC
3. Explain and use the pH scale
4. Given concentrations of H+ & OH-,
calculate pH
5. Given pH, calculate concentrations of
H+ & OH-
Self-Ionization of Water
Two water molecules produce a hydronium
ion and a hydroxide ion by transfer of a
proton

H2O(l)  H2O(l)  H3O (aq)  OH (aq)
-
Conductivity shows concentrations of
H3O+ and OH- are 1.0 x 10-7 mol/L at 25oC
Kw = [H3O+][OH-] = 1.0 x 10-14 M2
Self-Ionization of Water
Neutral, Acidic, and Basic
Solutions
Neutral solutions: have equal [H+] and [OH-]
Acidic Solutions: have greater [H+] than
[OH-]
Basic Solutions (alkaline): have greater
[OH-] than [H+]
Calculating [H+] and [OH-]
Kw = [H+] [OH-]
If given the concentration of one ion, the
concentration of the other can be
calculated
Example
A 1.0 x 10-4 M solution of HNO3 has been
prepared for a lab experiment. Calculate
+
[H3O ] and [OH ]
HNO3 is a strong acid so assume 100%
dissociation
-
HNO3(l) + H2O(l)  H3O (aq) + NO3 (aq)
+
Solution
+
-
Given: Kw = [H ] [OH ]
1.0 x 10-4 M HNO3
Find concentration of H+:
mol HNO 3 1 mol H 3O  mol H 3O 


 molarity of H 3O 
L solution 1 mol HNO 3 L solution
Solution
Solve algebraically:
Kw
[OH ] 

[H ]
-
Substitute:
-14
1
.
0
x
10
[OH - ] 
-4
1.0 x 10
[OH ]  1.0 x 10
-
-10
M
pH Scale
• pH – pouvior hydrogène or “hydrogen
power”
• pH is defined as the negative logarithm of
the hydronium ion concentration [H3O+]
• To calculate pH
pH = -log[H3O+]
or
pH = -log[H+]
pH Scale
• To calculate pOH
pOH = -log[OH-]
• At 25oC, [H+] = [OH-] = 1.0 x 10-7 = pH =7
+
• If [H ] > [OH ] the solution is acidic
• If [OH-] > [H+] the solution is basic
Examples:
1. What is the pH of 1.0 x 10-3 M NaOH
solution?
2. What is the pH of a solution if the [H3O+] is
3.4 x 10-5 M?
3. Determine the [H3O+] of an aqueous
solution that has a pH of 4.0.
4. The pH of a solution is found to be 7.52.
a) Find [H3O+]
b) Find [OH-]
c) Is the solution acidic or basic?
16-2 Objectives
• Describe how a pH indicator works
• Explain an acid-base titration
• Calculate molarity of a solution from
titration data
Indicators and pH Meters
Acid-Base Indicator – compounds whose
colors are sensitive to pH

HIn  H  In

In acidic solutions, In- ions act as B-L bases
In basic solutions, OH- react with H+ from
indicator
Indicators
Transition Interval – the pH range over
which an indicator changes color
Universal Indicators – combination of
different indicators
pH meters – determine pH by measuring
the voltage difference between two
electrodes
Titrations
The controlled addition of a measurement of the
amount of a solution of known concentration
required to react completely with a measured
amount of a solution of an unknown concentration
Equivalence Point – the point in a titration where [H+]
and [OH-] are present in chemically equivalent
amounts
End Point – the point of the titration at which the
indicator changes color
Molarity and Titration
Standard Solution – the solution that
contains the precisely known
concentration of a solute
Primary Standard – a highly purified solid
compound used to check the
concentration of the known solution in a
titration
Calculating Molarity from
Titration Data
1. Write balanced chemical equation and
determine chemical equivalents
2. Determine moles of acid or base from
the known solution used during the
titration
3. Determine moles of solute of the
unknown used during the titration
4. Determine molarity
Example: in a titration, 27.4 mL
of 0.0154 M NaOH is added to
a 20.0 mL sample of HCl
solution of unknown
concentration…find the
molarity of the acid solution
Solution
A 15.5 mL sample of 0.215 M KOH
solution required 21.2 mL of aqueous
acetic acid solution in a titration
experiment. Calculate the molarity of
the acetic acid solution.
Solution
By titration, 17.6 mL of aqueous H2SO4
neutralized 27.4 mL of 0.0165 M LiOH
solution. What was the molarity of
the aqueous acid solution?
Solution