Activity 151-15
Titrations & pH
Directions: This Guided Learning Activity (GLA) focuses on chemical calculations related to acids, bases
and pH. Part A gives basic information about acids and bases, and describes the ionization of water. Part
B discusses KW and the pH scale. Part C uses stoichiometry to solve titration problems. The worksheet is
accompanied by instructional videos. See http://www.canyons.edu/Departments/CHEM/GLA/ for
additional materials.
Part A – The Basics of Acids & Bases
Early in the class, you probably learned that an acid is identifiable because the chemical formula is
written with hydrogen as the first element. Acids are a class of chemicals that are able to ionize in
aqueous solutions to produce hydronium, H3O+, ions. For example, hydrochloric acid, or HCl, is a
covalent compound. Yet when it is placed in an aqueous solution, the hydrogen from the HCl binds to a
water molecule to form a hydronium ion, leaving the chloride ion in solution.
𝑯𝑯𝑯𝑯𝑯𝑯 (𝒈𝒈) + 𝑯𝑯𝟐𝟐 𝑶𝑶(𝒍𝒍) → 𝑯𝑯𝟑𝟑 𝑶𝑶+ (𝒂𝒂𝒂𝒂) + 𝑪𝑪𝒍𝒍− (𝒂𝒂𝒂𝒂)
All acids will ionize to some degree to form H3O+ ions in solution. Because of this, all acids share some
characteristics. For example acids are typically sour, can dissolve metals, and turn litmus paper red.
Bases, on the other hand, are compounds that increase the hydroxide, or OH-, concentration in solution. In
Chemistry 151, the bases you will encounter are soluble or slightly soluble hydroxide salts. Bases have a
bitter taste, a slippery soap-like feel, and will turn litmus blue.
The reason acids and bases are chemically important is because the presence of an acid or base will affect
the concentration of H3O+ and OH- in solution. In pure water, a very small fraction of water molecules
reacts to form both of these two ions:
𝑯𝑯𝟐𝟐 𝑶𝑶(𝒍𝒍) + 𝑯𝑯𝟐𝟐 𝑶𝑶(𝒍𝒍) ⇌ 𝑯𝑯𝟑𝟑 𝑶𝑶+ (𝒂𝒂𝒂𝒂) + 𝑶𝑶𝑯𝑯− (𝒂𝒂𝒂𝒂)
Adding an outside source of H3O+ of OH- will affect the fraction of water molecules that are ionized.
Whenever the amount of H3O+ is increased, the amount of OH- in the solution will decrease. When the
amount of OH- is increased, the amount of H3O+ decreases. This occurs because when the ions encounter
each other, they react to form water.
Often, the hydronium ion is written as simply H+. This communicates the reactive portion of the ion (the
ionized hydrogen from the acid). In general, H+ and H3O+ can be used interchangeably. But you should
know that a hydrogen ion (H+) cannot exist in isolation, is always ‘carried’ by another substance –
typically another water molecule. In this GLA, the hydronium ion is written only as H3O+.
Chemistry Guided Learning Activities
Activity 151 – 15
College of the Canyons
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Practice:
Classify each of the following as an acid, a base, or neither; then determine whether the compound will
release H3O+, OH- or neither into solution.
HCl is
an acid
and produces
H3O+
in solution.
SrCl2
neither
and produces
n/a
in solution.
HC2H3O2
__________
and produces
_________
in solution.
KOH
__________
and produces
_________
in solution.
NaF
__________
and produces
_________
in solution.
Ba(OH)2
__________
and produces
_________
in solution.
Part B – KW and the pH Scale
As mentioned above, adding an acid or a base to water will change both the amount of H3O+ and the
amount of OH- in the solution. Usually, the amount of H3O+ and OH- is expressed in molarity because
acids and bases exist in aqueous solutions. (See GLA 151-14 for a review of molarity and concentration
units). The two concentrations are related by the equation:
[𝑯𝑯𝟑𝟑 𝑶𝑶+ ] ∗ [𝑶𝑶𝑯𝑯− ] = 𝟏𝟏𝟏𝟏𝟏𝟏𝟏𝟏−𝟏𝟏𝟏𝟏
Where [H3O+] is the molar concentration of hydronium ions, [OH-] is the molar concentration of
hydroxide ions, and 1x10-14 is a constant called “KW”, or the ionization constant for water.
Example #1. A basic solution contains 0.0004 M hydroxide ions. What is the concentration of H3O+ in
this solution?
Because only a very small fraction of water ionizes, typical concentrations of H3O+ and OH- are quite low
(~1x10-7 M). This characteristic has led to the development of the pH scale. The pH scale is defined based
on the molar concentration of H3O+ in solution:
This can also be expressed as:
Chemistry Guided Learning Activities
Activity 151 – 15
𝒑𝒑𝒑𝒑 = −𝐥𝐥𝐥𝐥𝐥𝐥[𝑯𝑯𝟑𝟑 𝑶𝑶+ ]
[𝑯𝑯𝟑𝟑 𝑶𝑶+ ] = 𝟏𝟏𝟏𝟏−𝒑𝒑𝒑𝒑
College of the Canyons
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According to this scale, acidic solutions (which have a high [H3O+]) will have a low pH (below 7) and
basic solutions (which have a high [OH-]) will have a high pH (above 7). Solutions that have a pH equal
to 7 are considered neutral and have [H3O+] = [OH-]. There is no physical limit to the pH scale, but it is
generally drawn with boundaries around 0 and 14. Keep in mind that if we are given pH, we can
determine [OH-] by using the Kw mentioned earlier.
Example #2. What is the pH of a solution with a H3O+ concentration of 1x10-4 M? Is the solution acidic,
basic, or neutral?
Part C – Titrations
Titrations are a specific application of the stoichiometry concepts already discussed in GLA 151-9.
Specifically, an acid-base titration is a double replacement reaction between an acid and a base. These
reactions will always form a salt and water. (Refer to GLA-6 – Predicting Products in Chemical Reactions
for additional guidance.)
Because titrations are performed in aqueous solutions, generally the amount of reactant is given in terms
of molarity, a concentration. Remember, a 1.0 molar solution contains 1.0 mole of solute in 1.0 liter of
solution. A molar concentration can easily be written as a conversion factor. (Refer to GLA 151-14 –
Units of Concentration for additional guidance.)
0.3 𝑀𝑀 𝑁𝑁𝑁𝑁𝑁𝑁𝑁𝑁 =
0.3 𝑚𝑚𝑚𝑚𝑚𝑚 𝑁𝑁𝑁𝑁𝑁𝑁𝑁𝑁
0.3 𝑚𝑚𝑚𝑚𝑚𝑚 𝑁𝑁𝑁𝑁𝑁𝑁𝑁𝑁
=
1 𝐿𝐿 𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠
1000 𝑚𝑚𝑚𝑚 𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠
During a titration, a solution with a known concentration (a titrant) is added slowly to another solution
with an unknown concentration (an analyte). Titrations utilize indicators that undergo some change,
typically a color change, when the amount of titrant added is enough to react completely with the analyte.
This is termed the equivalence point. The equivalence point is where you see the most dramatic pH
change with even small additions of titrant. Because both the concentration of the titrant and the amount
of titrant added are known, the amount of analyte present can be found. (Refer to GLA 151-9 –
Introduction to Stoichiometry for additional guidance.)
Chemistry Guided Learning Activities
Activity 151 – 15
College of the Canyons
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Example #3. Label the analyte, titrant, indicator, and equivalence point in the following diagram.
14
12
10
pH
8
6
4
2
0
0
5
10
mL Titrant Added
15
Example #4. 50.0 mL of a hydrobromic acid solution is titrated with 28.34 mL of 0.09683 M NaOH.
How many moles of hydrobromic acid were present in the 50.0 mL solution?
Solution:
𝑯𝑯𝑯𝑯𝑯𝑯(𝒂𝒂𝒂𝒂) + 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵(𝒂𝒂𝒂𝒂) → 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵(𝒂𝒂𝒂𝒂) + 𝑯𝑯𝟐𝟐 𝑶𝑶(𝒍𝒍)
𝟎𝟎. 𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎 𝒎𝒎𝒎𝒎𝒎𝒎 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵
𝟏𝟏 𝒎𝒎𝒎𝒎𝒎𝒎 𝑯𝑯𝑯𝑯𝑯𝑯
𝟏𝟏 𝑳𝑳 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵
��
��
�
𝟏𝟏 𝑳𝑳 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵
𝟏𝟏 𝒎𝒎𝒎𝒎𝒎𝒎 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵
𝟏𝟏𝟏𝟏𝟏𝟏𝟏𝟏 𝒎𝒎𝒎𝒎 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵
= 𝟎𝟎. 𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎𝟎 𝒎𝒎𝒎𝒎𝒎𝒎 𝑯𝑯𝑯𝑯𝑯𝑯
(𝟐𝟐𝟐𝟐. 𝟑𝟑𝟑𝟑 𝒎𝒎𝒎𝒎 𝑵𝑵𝑵𝑵𝑵𝑵𝑵𝑵) �
Example #4b. What was the concentration of the initial hydrobromic acid solution?
Example #5. How many moles of hydrochloric acid are needed to react completely with a 75.0 mL
solution of 0.0455 M barium hydroxide?
Chemistry Guided Learning Activities
Activity 151 – 15
College of the Canyons
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Part D – Extra Practice
1. What is the [OH-] in a solution with a pH of 9.4? Is the solution acidic, basic, or neutral?
2. What is the pH of 12 M hydrochloric acid? Is the solution acidic, basic or neutral? (*Hint: HCl is a
strong acid, so it ionizes completely, producing 12 M H3O+ ions.)
3. Complete the following table:
[H3O+]
[OH-]
pH
Acidic, basic or
neutral?
1.4 x 10-5
10.5
-11
2.3 x 10
Neutral
4. Determine the pH, pOH, [H3O+], or [OH-] for the following scenarios.
a. If [H3O+] = 0.10 M,
pH = ____________
b. If [H3O+] = 0.10 mM,
[OH-] = ____________
c. If [OH-] = 10 μM,
pH = ____________
d. If [OH-] = 0.00025 M,
[H3O+] = ____________
e. If [OH-] = 0.00025 M,
pH = ____________
f. If [OH-] = 0.00025 M,
pOH = ____________
5. 15.0 mL of H3PO4 is titrated with 0.0100 M NaOH. 47.5 mL of the NaOH solution is needed to reach
the equivalence point. What is the concentration of H3PO4 in the original solution?
6. 25.0 mL sample of acetic acid (unknown concentration) was titrated with 32.98 mL of 0.112 M KOH.
What is the concentration of the acetic acid?
7. How much 0.055 M Ba(OH)2 is required to reach the equivalence point when titrating 16.7 mL of a
solution of 2.00% (m/v) HC2H3O2?
Chemistry Guided Learning Activities
Activity 151 – 15
College of the Canyons
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