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QUANTITATIVE TITRATION
(Revised 5/29/96)
A “titration” is an experimental process for determining the volume of a solution needed
to react with a given mass or volume of a sample. You will use this process to study
quantitatively the reaction between an acid and a base. The reaction of the hydrogen
ion (H+) of an acid with the hydroxide ion (OH-) of a base to form water is a common
reaction that will be used in this lab.
In an acid-base titration, the important information to obtain is the equivalence point. If
there are a given number of moles of acid in the titration flask, the equivalence point is
reached when the same number of moles of base have been added from the buret.
Often the pH of the solution will change dramatically at the equivalence point. An acidbase indicator works by changing color over a given pH range. The best choice for an
indicator will drastically change color near the equivalence point of the titration. The
drastic and rapid change in the indicator’s color is due to the rapid change of the pH of
the solution near the equivalence point.
In a potentiometric acid-base titration, an indicator is not necessary. A pH meter is used
to measure the pH as base is added in small increments (called aliquots) to an acid
solution. A graph is then made with the pH along the vertical axis and the volume of
added base along the horizontal axis. From this graph the equivalence point can be
determined.
OBJECTIVES:
1. To perform a potentiometric titration of an acidic solution of known molarity and
calculate the molarity of the base.
2. To perform a potentiometric titration of a solution containing a known amount of an
unknown, monoprotic solid acid, using the base of a known molarity as the titrant.
The molar mass of the acid will then be calculated.
3. To perform a potentiometric titration of a sample of vinegar, using the same base as
the titrant. The mass percent of acetic acid in vinegar will be calculated.
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IMPORTANT FACTS ABOUT THE pH PROBES
1. The pH probes are fragile and they may break if they get hit by the stirring bar.
Always adjust the pH probe so that it will not come into contact with the stir bar, as
illustrated in figure one. position and keep the stirring motor on the lowest setting.
2. After using the pH probes they should be rinsed well with distilled water and then
gently blotted with a KIMWIPE. DO NOT USE A PAPER TOWEL ON THE pH
PROBE!
3. The pH probe must not be allowed to dry out. The probe is to be stored in its
protective cap, or held in place in distilled water. Do not remove the protective cap
until you are ready to use the pH meter and replace the cap when you are done.
When storing the pH probes, put a few drops of pH = 4 or 7 buffer in the cap before
replacing the cap.
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EXPERIMENT A: DETERMINATION OF THE MOLARITY OF THE BASE SOLUTION
1. Obtain about 100 mL of the base solution in a clean, dry beaker. This beaker should
be labeled. Never pour any solution back into this beaker. Once the solution has
been poured into the buret, it should be discarded into the sink or the waste
container.
2. Rinse the buret with distilled water. Then use a small amount of the base solution to
rinse the buret three times. This is done by pouring about 10 mL of the solution into
the buret, letting some of the solution flow through the tip, and then pouring the rest
of the solution out the top of the buret, rotating the buret as you pour.
3. Fill the buret with the base solution to some point higher than the markings and then
slowly let the solution out into a waste container until the bottom of the meniscus is
on the 0.00 mL.
CALIBRATING THE METER AND ELECTRODE
* see supplemental pH meter instructions
TAKING pH MEASUREMENTS
The pH of the samples will be measured by immersing the electrode in the
solution. Allow the pH meter time to stabilize before recording the pH
value.
4. Using a 10 mL pipet or graduated cylinder, transfer 20 mL of 0.15 M HCL into a 400
mL beaker. If you use a graducated cylinder, use a medicine dropper to raise the
level of the meniscus to the 10.0 ml mark. With a pipet, adjust the level to the 10.00
ml line.
5. Add 50 mL of distilled water. Carefully drop a magnetic stirring bar into the beaker
and set the beaker on the stirring plate as shown in figure one. Adjust the pH probe
so that it will not be hit by the stirring bar. Turn the stirring motor onto the lowest
setting. If the solution level is to low, add enough distilled water to keep the stir bar
from hitting the probe.
6. Record the pH of the solution. The add 1.0 mL of the NaOH solution from the buret.
When the pH meter has stabilized, record the pH. Add another 1.0 mL and record
the pH. Continue adding NaOH in 1.0 mL increments until you have obtained a pH
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reading greater than 12 or when the pH reading seem to be leveling off.
7. Rinse the pH probe with distilled water and blot with a Kimwipe. Place three or four
drops of pH 4 or 7 buffer in the wetting cap and replace it on the electrode tip.
CALCULATIONS:
1. Make a graph of the pH vs. mL of NaOH added. The pH should be on the vertical
axis and the mL of NaOH should be on the horizontal axis. The graph should be of
such a size that 1 mL is represented by 1 square on the graph and the pH scale is
reasonably spread out.
2. There should be a region on your graph where the slope is very steep. Determine
the midpoint of this region. This is the equivalence point. Show on the graph the mL
of NaOH added at the equivalence point.
3. Calculate the number of moles of H+ in the beaker (HCL) . At the equivalence point
this will also be the number of moles of OH- from the NaOH.
4. Calculate the molarity of the base by dividing the number of moles of OH- by the mL
of NaOH at the equivalence point.
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EXPERIMENT B: DETERMINATION OF THE MOLAR MASS OF AN ACID
1. Obtain a solid unknown acid and record the number on the vial. Find the mass of the
vial containing the acid and record it in your data table. Empty the contents of the
vial into a 400 mL beaker. Find the mass of the empty vial and record it in your data
table.
2. Obtain about 100 mL of the base solution in a clean, dry beaker. This beaker should
be labeled. Never pour any solution back into this beaker. Once the solution has
been poured into the buret, it should be discarded into the sink or the waste
container.
3. Rinse the buret with distilled water. Then use a small amount of the base solution to
rinse the buret three times. (This is done by pouring about 10 mL of the solution into
the buret, letting some of the solution flow through the tip, and then pouring the rest
of the solution out the top of the buret, rotating the buret as you pour.)
4. Fill the buret with the base solution to some point higher than the markings and then
slowly let the solution out into a waste container until the bottom of the meniscus is
on the 0.00 mL line.
CALIBRATING THE METER AND ELECTRODE
*see supplemental pH meter instructions
TITRATION PROCEDURE
5. Add 50 mL of distilled water to the beaker containing the solid acid, and carefully
place the magnetic stirring bar in the beaker. Carefully drop a magnetic stirring bar
into the beaker and set the beaker in the magnetic stirring motor as shown. Adjust
the pH probe (figure one) so that it will not be hit by the stirring bar. Turn the stirring
motor onto the lowest setting.
6. Record the pH of the solution. Then add 1.0 mL of the NaOH solution from the buret.
When the pH meter has stabilized, record the pH. Add another 1.0 mL and record
the pH. Continue adding NaOH in 1.0 mL increments until you have obtained a pH
reading greater than 12 or when the pH reading seem to be leveling off.
7. Rinse the pH probe with distilled water and blot with a Kimwipe. Place three or four
drops of pH=4 or 7 buffer into the wetting cap and replace it on the electrode tip.
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CALCULATIONS:
1.
Make a graph of the pH vs. mL of NaOH added. The pH should be on the vertical
axis
and the mL of NaOH should be on the horizontal axis. The graph should be of
such a
size that 1 mL is represented by one square on the graph and the pH scale is
reasonably
spread out.
2.
the
There should be a region on your graph where the slope is very steep. Determine
midpoint of this region. This is the equivalence point. Show on the graph the mL of
NaOH added at the equivalence point.
3.
Calculate the number of moles of OH- added to the beaker at the equivalence
point. This
will also be the number of moles of acid.
4.
Calculate the mass of the acid.
5.
Calculate the molar mass of the acid by dividing the mass of the acid by the
number of
moles.
6.
Obtain the actual molar mass from your teacher and calculate the percent
difference.
% difference =
your value − actual value
* 100
actual value
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EXPERIMENT C: MASS PERCENTAGE OF ACETIC ACID IN VINEGAR
1. Measure out 5.0 mL of vinegar. Place a 400 mL beaker on the balance and TARE it
to 0.00 g. Add the 5.0 mL of vinegar and record the mass in your data table.
2. Obtain about 100 mL of the base solution in a clean, dry beaker. This beaker should
be labeled. Never pour any solution back into this beaker. Once the solution has
been poured into the buret, it should be discarded into the sink or the waste
container.
3. Rinse the buret with distilled water. Then use a small amount of the base solution to
rinse the buret three times. (This is done by pouring about 10 mL of the solution into
the buret, letting some of the solution flow through the tip, and then pouring the rest
of the solution out the top of the buret, rotating the buret as you pour.)
4. Fill the buret with the base solution to some point higher than the markings and then
slowly let the solution out into a waste container until the bottom of the meniscus is
on the 0.00 mL line.
CALIBRATING THE METER AND ELECTRODE
*see supplemental pH meter instructions
TITRATION PROCEDURE
5. Add 50 mL of distilled water to the beaker containing the vinegar, and carefully place
the magnetic stirring bar in the beaker. Carefully drop a magnetic stirring bar into the
beaker and set the beaker in the magnetic stirring motor as shown. Adjust the pH
probe (figure one) so that it will not be hit by the stirring bar. Turn the stirring motor
onto the lowest setting.
6. Record the pH of the solution. Then add 1.0 mL of the NaOH solution from the buret.
When the pH meter has stabilized, record the pH. Add another 1.0 mL and record
the pH. Continue adding NaOH in 1.0 mL increments until you have obtained a pH
reading greater than 12 or when the pH reading seem to be leveling off.
7. Rinse the pH probe with distilled water and blot with a Kimwipe. Place three or four
drops of pH=4 or 7 buffer into the wetting cap and replace it on the electrode tip.
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CALCULATIONS:
1. Make a graph of the pH vs. mL of NaOH added. The pH should be on the vertical
axis and the mL of NaOH should be on the horizontal axis. The graph should be of
such a size that 1.0mL is represented by 1 square on the graph and the pH scale is
reasonably spread out.
2. There should be a region on your graph where the slope is very steep. Determine
the midpoint of this region. This is the equivalence point. Show on the graph the mL
of NaOH added at the equivalence point.
3. Calculate the number of moles of OH- added to the beaker at the equivalence point.
This will also be the number of moles of acetic acid.
4. The molecular formula for acetic acid is HC2H3O2. Calculate the mass of the acid
that was titrated.
5. Calculate the mass percent of acetic acid in vinegar by dividing the mass of acetic
acid titrated by the mass of sample of vinegar.
6. Record the mass percent of acetic acid given on the vinegar bottle.
7. Calculate the percent difference.
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QUANTITATIVE TITRATION: TEACHER’S GUIDE
CLASSROOM USAGE & TIME
This is a simple acid-base titration experiment. The students can develop
skills in manipulating the buret’s stopcock, finding the equivalence point from a graph of
the titrations, and utilizing the mole concept. The three experiments including the
calculations and graphs can be completed in three 50 minute class periods.
PREVIOUS KNOWLEDGE
Neutralization Reactions
Titration
Acid-Base Ionization
pH
Molarity
Mole Concept
Ionic & Covalent Bonding
SAFTEY AND DISPOSAL
The instructor(s) and the students need to wear goggles and lab aprons during the
lab. Standard lab safety procedures should be followed. All waste reagents may be
washed down the drain with water.
PREPARATION: EQUIPMENT
12 pH meters & probes
12 stir plates
12 50mL graduated cylinders
graph paper
12 10mL graduated cyllinders
12 25mL burets
wash bottles
12 10ml or 20ml pipets & bulbs (if available)
PREPARATION: REAGENTS
2 L distilled water
2 Buffer Solutions: pH 7 & pH 10
100 mL Vinegar
12 numbered vials @ ~0.5 g each
(if multiple solid acids are used)
4 L NaOH solution: 0.160 - 0.180 M
0.5 L HCl solution: 0.150 M
Solid Monoprotic Acid (i.e. KHP)
PREPARATION NOTES
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1. Buffer Solutions: If the pH meter allows calibration for only two buffers, use the pH=4
& pH=7 buffers. If the pH meter allows calibration for more than two buffers, use the
pH=10 buffer in addition to the first two.
2. The NaOH solution’s concentration can be approximately 0.160 M to 0.180 M. If you
make the concentration the same as the HCl solution’s concentration, the students may
make the connection that there is no difference betwenn the solutions and that ther are
interchangable.
3. The HCL solution has to be made up as quatitaively as possible. The following is a
procedure for making the standard 0.150 aAM HCl solution as taken from Parry, Dietz,
Tellefsen, & Steiner, Chemistry: Experimental Foundations, 2nd ed. In order to make a
smaller amount that the procedure describes, the amounts of water and concentrated
HCl can be cut in half, but no more.
“The standard 0.150 M acid should be made up in a large batch to insure uniformity.
Measure carefully 128 ml of concentrated (11.7M) HCl (fresh bottle). Place in a 1-liter
volumetric flask about half-filled with water. After mixing, fill to the 1-liter mark with
distilled water. Empty this into a large container, and add 9 more liters of distilled water.
Acid diluted in this manner need not be standardized. Five gallon water jugs and
convenient for mixing and storing this large volume.”
4. Here is a listing of possible solid monoprotic acids.
Acid
Formula
Molecular
Weight
Grams*
potassium hydrogen
sulfate
KHSO4
136.2
0.8
potassium hydrogen
tartrate
KHC4H4O6
188.2
1.1
potasium hydrogen
phthalate
KHC8H4O4
(KHP)
204.2
1.2
potassium hydrogen oxlate
KHC2O4
128.1
0.8
* amount needed to neutralize 50ml of 0.12 M NaOH
Lab Written By: Pru Phillips
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SAMPLE DATA
Experiment A: Determining the Molarity of a Base.
mL NaOH
pH
mL NaOH
pH
mL NaOH
pH
0
1.06
11
1.35
22
11.80
1
1.04
12
1.41
23
11.94
2
1306
13
1.48
24
12.03
3
1.07
14
1.56
25
12.10
4
1.09
15
1.69
26
12.18
5
1.11
16
1.81
27
12.21
6
1.14
17
2.01
28
12.25
7
1.16
18
2.28
29
12.30
8
1.19
19
9.20
30
12.32
9
1.24
20
11.24
(3) 0.0200 L HCl *
moles HCl
0150
.
= 0.00300 moles HCl titrated
1L
0.00300 moles HCl titrated = 0.00300 moles NaOH delivered
(4)
10
0.00300 moles NaOH
= 0.161 M NaOH
0.0186 L
1.30
21
11.60
(1) & (2) See next page
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Experiment B: Determination of the Molar Mass of an Acid
mL NaOH
pH
mL NaOH
pH
mL NaOH
pH
0
3.85
11
4.88
22
11.69
1
3.93
12
4.97
23
11.90
2
4.08
13
5.06
24
12.09
3
4.19
14
5.15
25
12.22
4
4.29
15
5.25
26
12.31
5
4.39
16
5.38
27
12.39
6
4.47
17
5.53
28
12.45
7
4.55
18
5.74
29
12.51
8
4.63
19
6.06
30
12.57
9
4.71
20
6.95
10
4.80
21
11.35
(1) & (2) See next page
(3) 0.0204 L NaOH *
0.161 moles NaOH
= 0.00328 moles NaOH delivered
1L
0.00328 moles NaOH delivered = 0.00328 moles acid titrated
(4) & (5)
(6)
0.660g acid
= 201 g / mole (calculated molar mass of the acid)
0.00328 moles acid
201g / mole - 204 g / mole
204 g / mole
∗100 = 1.47% difference
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Experiment C: Mass Percent of Acetic Acid in Vinegar
mL NaOH
pH
mL NaOH
pH
mL NaOH
pH
0
2.58
12
4.23
24
5.17
1
2.98
13
4.30
25
5.32
2
3.27
14
4.36
26
5.49
3
3.43
15
4.42
27
5.76
4
3.56
16
4.49
28
6.34
5
3.70
17
4.56
29
10.94
6
3.78
18
4.63
30
11.67
7
3.88
19
4.71
31
11.85
8
3.97
20
4.79
32
12.02
9
4.03
21
4.87
33
12.13
10
4.10
22
4.95
34
12.22
11
4.17
23
5.04
35
12.29
(1) & (2) See next page
(3) 0.0285 L NaOH *
0161
moles NaOH
.
= 0.00459 moles NaOH delivered
1L
0.00459 moles NaOH delivered = 0.00459 moles acetic acid titrated
(4) 0.00459 moles of acetic acid *
(5)
(6)
60.0 g acetic acid
= 0.275 g acetic acid tritrate
1 mole acetic acid
0.275 g acetic acid
* 100 = 5.01% acetic acid
5.50 g sample vinegar
5.01 % - 5.00%
5.00 %
∗100 = 0.200% difference
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