5.3.1 Neutralization reactions
5.3.2 Titration Reactions
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What happens when an acid such as HCl is
mixed with a base such as NaOH:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
When an acid and a base are combined,
water and a salt are the products.
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Salts are ionic compounds containing a
positive ion other than H+ and a negative ion
other than the hydroxide ion, OH-.
Double displacement reactions of this type
are called neutralization reactions.
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We can write an expanded version of this
equation, with aqueous substances written in
their longer form:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl(aq) + H2O(l)
Removing the spectator ions we get the net
ionic equation:
H+(aq) + OH-(aq) → H2O(l)
When a strong acid and a strong base are
combined in the proper amounts ([H+] equals
[OH-]) - a neutral solution results (pH = 7).
 But, salt solutions do not always have a pH of 7.
 Through a process known as hydrolysis, the ions
produced may react with water molecules to
produce a solution that is slightly acidic or basic.
 Generally if a strong acid is mixed with a weak
base there the resulting solution will be slightly
acidic; if a strong base is mixed with a weak acid
the solution will be slightly basic.
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Acid-base titrations are lab procedures used
to determine the concentration of a solution
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During an acid-base titration, an acid with a
known concentration (a standard solution) is
slowly added to a base with an unknown
concentration (or vice versa). A few drops of
indicator solution are added to the base.
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The indicator will change colour, when the
base has been neutralized (when [H+] =
[OH-]).
At that point - called the equivalence point
or end point - the titration is stopped. By
knowing the volumes of acid and base used,
and the concentration of the standard
solution, calculations allow us to determine
the concentration of the other solution.
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It is important to accurately measure volumes
when doing titrations. The instrument you
would use is called a burette (or buret).
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1.Rinse the burettes several times with the
acid and base solutions, then fill them with
the appropriate solution.
2.Using the burette, accurately measure a
volume of the base into an Erlenmeyer flask.
How much you use is not important, as long
as you know exactly how much is in the flask.
3.Add a suitable indicator such as
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4. The acid (with the known concentration - the
titrant) is then added to the base. Release the acid
from the burette quickly at first, then more slowly as
the endpoint is neared.
How do you know when you are reaching the
endpoint? The indicator will begin to show a change in
colour. Swirling the flask will cause the colour to
disappear.
ENDPOINT IS REACHED AS SOON AS THE
COLOUR CHANGE IS PERMANENT. ONE DROP
WILL DO IT - once the colour change has occurred,
stop adding acid.
If a pH meter is used instead of an indicator, endpoint
will be reached when there is a sudden change in pH.
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5. Record the volume of acid added to reach
endpoint. Calculations will allow us to
calculate the concentration of the base.
6. The titration should be repeated several
times and the results averaged. Often the
first trial is done quickly to get a general idea
of the volume of titrant required. Subsequent
trials are done more slowly.
Use titration information to calculate the
concentration of the unknown solution, you
must know the following information.
 The concentration of one of the solutions, the
acid for example (MA)
 The volume of acid used for the titration (VA)
 The volume of base used for the titration (VB)
 What you will calculate:
 The concentration of the other solution, the
base for example (MB)
MAVA = MBVB
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During a titration 75.8 mL of a 0.100 standard
solution of HCl is titrated to end point with
100.0 mL of a NaOH solution with an
unknown concentration. What is the
concentration of the NaOH solution?
MA= 0.100 M
MB = MB
VA =75.8 mL
VB =100.0 mL
Substitute in known values and solve for the
unknown:
MAVA = MBVB
(0.100)(75.8) = MB(100.0)
7.58 /100.0 = MB
0.0758 = MB
Answer: The concentration of the NaOH solution
is 0.0758 M or 7.58 ×10-2M
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A 20.0 mL solution of strontium hydroxide,
Sr(OH)2, is placed in a flask and a drop of
indicator is added. The solution turns colour
after 25.0 mL of a standard 0.0500M HCl
solution is added. What was the original
concentration of the Sr(OH)2 solution?
Write a balanced equation for the neutralization reaction:
2 HCl(aq) + Sr(OH)2 (aq) → SrCl2 (aq) + H2O(l)
Notice that 2 moles of the acid HCl are required to neutralize
1 mole of the base Sr(OH)2.
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Define the variables to be used and the known values:
MA = 0.050 M
MB = MB
VA = 25.0 mL
VB = 20.0 mL
To balance this out we need to modify our formula:
MAVA = 2 MBVB
Notice that the "2" is on the base side of our formula even
though it was in front of the acid side in the balanced
equation - it "switches places".
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Substitute in known values for the equation
and solve for the unknown:
MAVA = 2MBVB
(0.050)(25.0) = 2MB(20.0)
1.25 = 40.0MB
1.25/ 40.0 = MB
0.0312 = MB
Answer: The concentration of the
Sr(OH)2,solution is 3.12 ×10-2M.
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(You will not be required to know this section
for a test.)
What indicator is used depends on what type
of titration you are performing. An indicator
should be chosen that will change color when
enough of one substance (acid or base) has
been added to exactly use up the other
substance.
The three main types of acid-base titrations,
and suggested indicators, are:
Titration between . . . Indicator
1. strong acid
and strong base
any
2. strong acid
and weak base
methyl
orange
3. weak acid
Explanation
phenolphthalein
changes color
in the acidic
range(3.2-4.4)
changes
color and
strong base
in the basic
range (8.2 - 10.6)
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Lab 5.3.2- titrations
Assignment 5.3.2