Preparation and Standardization of NaOH
Objective
This lab focuses on the detection of ions using titration as an analysis tool. Because
NaOH is hygroscopic (i.e. it pulls moisture from the air) a precise concentration cannot
be achieved by simply weighing out the solid to make a solution. NaOH must be titrated
with a solution of a known concentration of acid to determine the exact concentration. In
this lab, the student will use prepare and standardize a NaOH solution that will be used
for future experiments.
Background
What is a Titration?
A titration is an analytical procedure used to determine the concentration of a sample by
reacting it with a standard solution. One type of titration uses a neutralization reaction, in
which an acid and a base react to produce a salt and water.
In equation 1, the acid is HCl (called hydrochloric acid) and the base is NaOH (called
sodium hydroxide). When the acid and base react, they form NaCl (sodium chloride),
which is also known as table salt. The titration proceeds until the equivalence point is
reached, where the number of moles of acid is equal to the number of moles of base.
This point is usually marked by observing a color change in an added indicator.
In a titration, the standard solution goes in a buret, which is a piece of glassware used to
measure the volume of solvent to approximately 0.01 mL of accuracy. The solution that
you are titrating goes in an Erlenmeyer flask, which should be large enough to
accommodate both your sample and the standard solution you are adding.
What is an Indicator and What is it Used For?
An indicator is any substance in solution that changes its color as it reacts with either an
acid or a base. Selecting the proper indicator is important because each indicator changes
its color over a particular range of pH values. Indicators are either weak acids or weak
bases. For example, phenolphthalein is a weak acid (which we will represent as HIn). In
aqueous solution, the phenolphthalein dissociates slightly, forming an equilibrium.
An equilibrium occurs when the amount of reactants and the amount of products are
constant. When a system is in equilibrium, it will stay there until something changes the
conditions. A famous French chemist, named Le Chatelier, developed a way to predict
how changes in equilibrium affect the system. Le Chatelier’s principle states that when
an equilibrium is disturbed by applying stress, the equilibrium will shift to relieve the
stress. In an acidic solution, there is an excess of H3O+ ions so the equilibrium will shift
to the left and favor the formation of HIn, thus we observe a clear solution. In basic
solution, there is an excess of OH- ions that react with the H3O+ ions to form water. This
shifts the equilibrium to the right because water is being formed and H3O+ ions are being
removed, thus we observe a pink solution. We can use this color change to determine
when the end of the titration has been achieved.
Table 1 lists common indicators and the pH range over which they change colors.
Table 1: Table of Indicators
thymol blue
Color change
interval (pH)
1.2 - 2.8
methyl orange
3.1 - 4.4
red
yellow
methyl red
4.4 - 6.2
red
yellow
chlorophenol red
5.4 - 6.8
yellow
Red
bromothymol blue
6.2 -7.6
yellow
Blue
phenol red
6.4 - 8.0
yellow
Red
thymol blue
8.0 - 9.6
yellow
Blue
Phenolphthalein
8.0 - 10.0
colorless
Red
alizarin yellow
10.0 -12.0
yellow
green
Indicator
Acid
Base
red
yellow
Standardizing a Sodium Hydroxide (NaOH) Solution
In a titration, it is critical to know the exact concentration of the titrant (the solution in the
buret which will be added to the unknown) in order to determine the concentration of the
solution being tested. We will standardize the ~0.1 M NaOH solution (the titrant) with
potassium hydrogen phthalate (KHP, KC8H4O4H) using phenolphthalein as the indicator.
KHP is a weak acid and reacts with base in the following way:
Pre-lab:
1. Read the procedure and set up a data table in your lab notebook to capture data.
2. Write a procedure for preparing the NaOH solution and review your plans with
your instructor
3. If the equivalence point of the KHP and NaOH titration is shown by the plot
below? Which indicator(s) could be used to detect the equivalence point of this
titration?
EEquiv
alence
Equivalence
Point
4. How many ml of 0.100 M NaOH should react with 0.5 g of KHP?
Procedure
Prepare the NaOH solution - Before Class (or after school)
Prepare 1 Liter of approximately 0.100 M NaOH. Use a volumetric flask, deionized
water and analytical balance to prepare the NaOH solution. Discuss your planned
procedure with your instructor prior to beginning.
Standardize the NaOH - During Class
1. Weigh ~0.4-0.6 g of dried KHP (MW = 204.23 g/mol) into an Erlenmeyer flask
and dissolve in 50-75 mL of distilled water. Record the amount of KHP and
water used.
2. Add 2-3 drops of indicator into the flask and titrate to the first permanent
appearance of pink. Near the endpoint, add the NaOH dropwise to determine the
total volume most accurately.
3. Rinse the previously cleaned buret with at least four 5-mL portions of the
approximately 0.100 M sodium hydroxide solution that you have prepared.
Discard each portion. Completely fill the buret with the solution and remove the
air from the tip by running out some of the liquid into an empty beaker- Make
sure that the lower pant of the meniscus is at the zero mark or slightly lower.
Allow the buret to stand for at least 30 s before reading the exact position of the
meniscus. Remove any hanging drops from the buret tip by touching it to the side
of the beaker used for the washings. Record the initial buret reading.
4. Slowly add the sodium hydroxide solution to one of your flasks of KHP solution
while gently swirling the contents of the flask. As the sodium hydroxide solution
is added, a pink color appears where the drops of the base come in contact with
the solution. This coloration disappears with swirling. As the equivalence point
(the point at which the amount of acid and base are equal) is approached, the color
disappears more slowly, at which time the sodium hydroxide should be added
drop by drop. It is most important that the flask be swirled constantly throughout
the entire titration. The equivalence point is reached when one drop of the
sodium hydroxide solution turns the entire solution in the flask from colorless to
pink. The solution should remain pink when it is swirled- Allow the titrated
solution to stand for at least 1 min so the buret will drain properly. Remove any
hanging drop from the buret tip by touching it to the side of the flask and wash
down the sides of the flask with a stream of water from the wash bottle. Record
the buret reading, Repeat this procedure at least twice or until three trials have
been completed that are precise within 1.0 percent.
Data
NaOH Solution
NaOH used __________________g
Calculate NaOH Molarity
Calculated molarity ________________M NaOH based on initial mass
Trial 1
g KHP
Mol KHP
NaOH Initial buret
reading (ml)
NaOH Final buret
Trial 2
Trial 3
reading (ml)
Volume NaOH (ml)
Molarity NaOH
Average NaOH M
Standard Deviation
Post Lab
1. Determine the average molarity and standard deviation of your NaOH.
2. The standardized valued will be used going forward for all future analytical work
with the NaOH solution prepared. Most likely, the standardized molarity is less
than the calculated molarity based on the mass NaOH used. Why is this the
expected case?
3. How would the molarity of NaOH change if not all of the KHP weighed out is
transferred into the flask.