This document describes a procedure to determine the concentration of acetic acid in vinegar through acid-base titration. Vinegar samples of known volume will be titrated with a sodium hydroxide solution of known molarity. The collected data on titrant volume will be used to calculate the moles of acetic acid present and from that, the molarity and percentage by mass of acetic acid in the vinegar samples. Students will perform the titration, record observations, calculate results, and report the average molarity and percentage of acetic acid determined.

1. Stoichiometry IV: Titration of Vinegar
PRE-LAB ASSIGNMENTS:
To be assigned by your lab instructor.
STUDENT LEARNING OUTCOMES:
• Learn how to use the stoichiometry of a reaction to relate moles of one thing to moles of
something else.
• Learn how to perform an acid-base titration.
• Be able to work stoichiometry problems involving molarity.
EXPERIMENTAL GOALS:
In this procedure, the molarity of acetic acid in vinegar and the percentage of acetic acid in
vinegar will be determined by a reaction with a solution of sodium hydroxide.
INTRODUCTION:
This lab is another procedure designed to introduce the concept of reaction stoichiometry.
In this procedure, the molarity of acetic acid in a solution of vinegar will be determined by
titration with an aqueous solution of sodium hydroxide.
A titration is a procedure that is often used for determining the concentration of a
solution. Most commonly, a standard solution of known concentration is reacted with a solution
of unknown concentration. By measuring the volume of standard solution that reacts with a
known volume of the unknown solution, the concentration of the solution can be calculated from
the reaction stoichiometry.
A common example of this process is an acid-base titration, in which an acid or base of
unknown concentration reacts with a base or acid of known concentration in a neutralization
reaction:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
In the example in Figure 1, base (OH-
) of known concentration from the buret is added to the
acid (H+
) of unknown concentration in the flask until the equivalence point is reached, when the
number of moles of OH-
added equals the number of moles of H+
from the acid originally
present.
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2. 86

3. H2O
OH-
OH-
H2O
OH-
H+
H
+
H+
H+
OH-
OH-
Base of known
concentration
OH-
OH-
OH-
OH-
OH-
OH-
equivalence point
all H+
has become H2O
some OH
-
added;
some H+
has become H2O
H2O
H2O
OH-
beginning of titration
no OH
-
added
H2O
Acid of unknown
concentration
OH-
OH-
H2O
H+
H+
H+
H+
H2O
Figure 1. A typical acid-base titration.
At the equivalence point, moles HCl = moles NaOH, so the concentration of the unknown HCl
solution can be calculated from the reaction stoichiometry:
HClmol
NaOHmol1
HClmol1
L
NaOHmol
NaOHL =××
HClM
soln.HClL
HClmol
=
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4. Burets are usually marked in increments of 0.1 mL, starting with 0.0 mL at the top and
50.0 mL at the bottom. The volume of the liquid in the buret can be measured very precisely to
the nearest 0.01 mL, which is enough to achieve a reasonable accuracy in most volumetric
analyses.
The end point of the titration, where we experimentally estimate the equivalence point to
be, is usually signaled by the color change of an acid-base indicator. The indicator is chosen in
such a way that the end point occurs as closely as possible to the equivalence point of the
titration. In many strong acid-strong base or weak acid-strong base titrations, phenolphthalein is
a good indicator. Phenolphthalein is an organic dye which is colorless in an acidic environment,
but pink in a basic environment (more precisely, it changes from colorless to pink over a pH
range of 8.0 to 9.6). As the sodium hydroxide titrant is added to the unknown acid solution, faint
swirls of pink may be observed, which disappear quickly as the base is neutralized. As the
titration nears the end point, the pink swirls take longer and longer to disappear. The end point is
reached when one excess drop of the titrant reacts with the phenolphthalein, producing a
permanent pink color which does not disappear when the solution is swirled.
When performing a titration, it is necessary to first determine the concentration of the
known solution as accurately as possible. This process is referred to as standardization. In this
experiment, the sodium hydroxide solution has already been standardized, and you will be
provided with the solution’s concentration.
In this titration, the concentration of acetic acid, HC2H3O2, in commercial vinegar will be
determined by titration against a standard solution of NaOH. The equation for this neutralization
reaction is:
HC2H3O2(aq) + NaOH(aq) → NaC2H3O2(aq) + H2O(l)
From the known concentration of the sodium hydroxide, and the volume of the solution that
emerges from the buret, the number of moles of NaOH can be determined, which allows the
number of moles of HC2H3O2 originally present and the concentration and weight percent of the
acid to be determined.
PROCEDURE:
1. Pour approximately 75 mL of the NaOH solution into a 100 mL beaker.
2. Rinse the buret well with tap water, then rinse twice with 5 mL portions of deionized water
and twice more with 5 mL portions of the NaOH solution.
3. Attach a buret clamp to a ring stand and place the buret in the clamp. Fill the buret with the
NaOH solution. It is important to ensure that the buret tip does not contain an air bubble, so
open the valve to allow some of the solution to run out of the buret until there are no more
bubbles of air running out of the buret tip.
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5. 4. Record the volume reading on your sodium hydroxide buret at the starting point for the
titration (2). The buret is calibrated in milliliters to the nearest 0.1 mL, so the volume
readings should be estimated to the nearest 0.01 mL. The 0.0-mL mark is at the top of the
buret and the 50.0-mL mark is near the bottom. To measure the volume delivered from a
buret, take the difference between the volume reading at the starting point of the titration, and
the final volume reading after the titration. It is not necessary that the starting point be
exactly zero, as long as it is accurately known.
5. Clean two 250 mL Erlenmeyer flasks, and rinse them thoroughly with deionized water. Dry
the outside of the flasks, and weigh each flask (10). (It is not necessary for the inside of the
flask to be dry.)
6. Select one of the available vinegar solutions, and record its identity on the report sheet.
Record the initial vinegar buret reading (6), deliver 10-12 mL of the vinegar into one of the
Erlenmeyer flasks, and record the final vinegar buret reading (5). Reweigh the flask and
record the mass (9). Repeat the process with the second Erlenmeyer flask; making sure to
use a different volume of the same vinegar solution as was used in the first flask.
7. Add 2 drops of phenolphthalein indicator to each of the samples of vinegar.
8. Position one of the Erlenmeyer flasks containing a vinegar sample under the buret. Make
sure the buret tip is inside the flask, so no drops of sodium hydroxide solution are lost. Place
a sheet of white paper underneath the flask to make the color changes more obvious.
9. Open the valve on the buret and begin to add sodium hydroxide solution to the vinegar
sample, while constantly swirling the flask. The addition may be rapid at first, but as soon as
you start to see swirls of pink color in the vinegar solution, slow the rate of addition down to
a fast dropwise rate. Swirl the solution continuously, and observe the color which develops
as each drop is added to the vinegar solution. As the titration gets closer and closer to the
end point, the swirls of pink color will take longer and longer to disappear — as long as the
color disappears between drops, however, you have not yet reached the end point. When the
pink color is no longer cleared by swirling, stop the fast drip. Make sure the pink color still
disappears,1
and continue the titration one drop at a time, swirling the solution after each
drop, until one drops turns the vinegar solution to a solid light pink color which does not
disappear on swirling. This is the end point of the titration. Record the volume on the buret
(1) and rinse the contents of the flask out in the sink.
10. Repeat the titration with the other vinegar sample.
11. Dispose of the sodium hydroxide solution in the beaker and buret in the sink. Rinse the buret
well with tap water, and then rinse it again a couple of times with deionized water (making
sure to run some water through the buret tip) before returning it to the stockroom. Rinse the
beakers and flasks with tap water and deionized water.
1
If the pink color persists at this point, you have overshot the end point, and need to repeat the titration with a fresh
sample of vinegar.
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6. CALCULATIONS:
1. Calculate the volume of the sodium hydroxide solution used in mL (3) and L (4). From the
concentration of sodium hydroxide you have been given, determine the number of moles of
sodium hydroxide used (12), and from the stoichiometry of the reaction, determine the moles
of HC2H3O2 present in the vinegar sample (13).
2. Calculate the volume of the vinegar solution used in mL (7) and L (8). From the number of
moles of HC2H3O2 present in the vinegar sample and the volume of the vinegar sample,
calculate the molarity of acetic acid in vinegar (14). Record the average molarity of the two
experiments (15).
3. Determine the mass of the vinegar sample (11). From the number of moles of HC2H3O2 in
the solution (13), determine the mass of HC2H3O2 in the vinegar sample (16). From the mass
of HC2H3O2 and the mass of the vinegar (14), determine the mass % of HC2H3O2 in the
vinegar sample (17). Record the average mass % of the two experiments (18).
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7. 91

8. LAB REPORT
Titration of Vinegar
Name ________________________________ Date _________ Report Grade ______
Partner ________________________________ Section _________
Vinegar sample used: ______
Concentration of sodium hydroxide solution: ____________
First
Determination
Second
Determination
1. Final reading of NaOH ____________ ____________
2. Initial reading of NaOH ____________ ____________
3. Volume of NaOH used (mL) ____________ ____________
4. Volume of NaOH used (L) ____________ ____________
5. Final reading of vinegar ____________ ____________
6. Initial reading of vinegar ____________ ____________
7. Volume of vinegar used (mL) ____________ ____________
8. Volume of vinegar used (L) ____________ ____________
9. Weight of flask + vinegar ____________ ____________
10
.
Weight of empty flask ____________ ____________
11
.
Weight of vinegar ____________ ____________
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9. 93

10. Calculation of Molarity and Mass Percent of
HC2H3O2 in Vinegar
First
Determination
Second
Determination
12
.
Moles of NaOH used (show calculations) ___________
_
___________
_
13
.
Moles of HC2H3O2 in vinegar (show calculations) ___________
_
___________
_
14
.
Molarity of HC2H3O2 in vinegar (show calculations) ___________
_
___________
_
15
.
Average molarity of HC2H3O2 in vinegar ____________
16
.
Mass of HC2H3O2 in vinegar (show calculations) ___________
_
___________
_
94

11. 17
.
Mass % HC2H3O2 in vinegar (show calculations) ___________
_
___________
_
18
.
Average Mass % of HC2H3O2 in vinegar ____________
95