Acid-Base Titration Analysis
Determination of the Acid Content in Vinegar
Introduction
Acids and bases react in aqueous solution to form a salt and water. This reaction, known
as a neutralization reaction, is actually one type of a double displacement reaction. Two examples
are given below.
HCl (aq)
hydrochloric
acid
H2SO4 (aq)
sulfuric
acid
+
NaOH (aq)
sodium hydroxide
(a base)

+
2 KOH (aq)
potassium hydroxide
(a base)

NaCl (aq)
sodium chloride
(a salt)
K2SO4 (aq)
potassium sulfate
(a salt)
+
HOH (l)
water
+
2 HOH (l)
water
Because there are many different acids and bases, neutralization reactions are very
common. If a neutralization reaction is run in the laboratory under controlled conditions, it can be
used to determine the concentration of the acid or the base in the reaction. This procedure,
known as an acid-base titration analysis, involves accurately measuring the volume of a base that
is required to neutralize a known volume of acid. In order to calculate the concentration of the acid
from the laboratory data, we must also know the concentration of the base used in the titration.
Thus to determine the concentration of an acid, we must have a base of known concentration and
the equipment to accurately measure the volume of the acid and base reacting. The concentration
of the base is given in moles of acid or base per liter of solution or molarity (M).
In this experiment we will determine the concentration of acetic acid
(HC2H3O2) in household vinegar. A known volume of vinegar will be
measured out with a volumetric pipette and titrated with a solution of
sodium hydroxide (NaOH) of known molarity. The NaOH solution will be
added from a buret until all of the acid in the vinegar has been neutralized.
Since we need to accurately measure the volume of NaOH used in the
titration, we must have a means of knowing when just enough of the base
(NaOH) has been added to react with all of the acid (HC 2H3O2), in the
vinegar. This point is called the end point. To detect the end point, an
acid-base indicator is added to the vinegar.
The indicator,
phenolphthalein, is colorless in acidic solutions but turns pink in basic
solution. The end point is when the vinegar turns very light pink and the
color persists even when you swirl the solution.
Meet your equipment
0
A buret allows you to accurately measure the amount of liquid you add
to a container.
10
The first point to notice about your buret is that the numbering is
“upside-down” compared to other measuring devices you have used.
Burets measure the amount of liquid delivered from them, not the
amount they contain. A full buret will read “0 mL” and an empty 50 mL
buret will read “50 mL.” You read the level in the buret from the bottom
of the meniscus, as you did with the graduated cylinder in experiment 2.
20
30
40
Before you start to pour anything into the buret, check that the stopcock
of the buret is closed. (Closed is horizontal. Open is vertical. If you try
to fill a buret with the stopcock in the vertical position, you’re likely to get
your shoes wet.)
To fill the buret, use a funnel. You might need a stool to stand on, too.
Put a waste beaker under the buret while you fill it, to catch any spills.
50
stopcock
If you notice that your buret is dribbling liquid even when the stopcock is
in the closed position, the stopcock may be assembled incorrectly.
Consult the instructor for help.
Equipment and Materials
Equipment:
Materials:
goggles
buret stand
50 mL burets (2)
125 mL Erlenmeyer flask
150 mL beakers (4)
100 mL graduated cylinder
wash bottle
distilled water
acetic acid (unknown concentration)
sodium hydroxide (1.00 M)
phenolphthalein indicator
Purpose:
To determine the unknown concentration (in units of mass %) of acetic acid in household vinegar.
Procedure
1. Obtain approximately 65 mL of 1.00 M NaOH in a clean, dry 150 mL beaker. Record the
molarity of the NaOH in the data table. If the solution has not already been prepared, then
measure out the mass of solid NaOH required to make a 100.0 mL solution of 1.00 M NaOH
(aq). Have your very wise instructor check your calculation of the mass of NaOH pellets you
will add to create this solution before you add distilled water.
2. Rinse one buret with a little of the NaOH solution, and pour the rinse into the sink.
3. Close the stopcock on the base buret and fill it to a level slightly above the 0.00 mL line.
Open the stopcock and allow a little of the NaOH to run out into a small waste beaker. This
helps to eliminate air bubbles in the tip of the buret.
4. Adjust the level of the NaOH in your buret so that the bottom of the meniscus is at or below
the 0.00 line. (If you need to remove some NaOH from the buret, let it flow out into the
waste beaker.)
5. Obtain approximately 65 mL of unknown concentration of vinegar in a clean, dry 150 mL
beaker.
6. Repeat steps 2 through 4 with the acetic acid (vinegar).
7. Transfer 20.00 mL of vinegar from the buret to a clean, dry 125 mL Erlenmeyer flask
8. Add 3 to 4 drops of phenolphthalein to this flask. The solution will be colorless.
9. Place your flask containing the 20.00 mL sample and the phenolphthalein under the tip of
the buret containing base. Carefully add the NaOH by turning the stopcock almost one half
turn. This allows one drop at a time to leave the buret. With experience, you may open the
stopcock and let more out at one time. Try to achieve an end point with the palest pink you
can see that persists for at least twenty seconds. Note: Some pink color may be observed
where the NaOH solution first contacts the vinegar, due to incomplete mixing. Swirling the
contents of the Erlenmeyer will remove the pink color if the end point has not yet been
reached.
10. If you pass the endpoint, then use some of the acid from the second buret to “back-titrate”.
This volume must be added to your recorded volume of “acid used”.
11. Once you have reached the end point, read and record the volume of NaOH that you have
added to your flask. Record this volume to the nearest 0.01 mL.
12. Subtract the initial volumes from the final volumes to calculate the titration volumes of acid
and base, the total volumes required for neutralization.
13. Perform a second titration with a fresh 20.00 mL sample of vinegar and phenolphthalein. If
the titration volumes for these two trials are not in good agreement, within 0.02 mL of each
other, do a third titration.
14. Calculate the average titration volumes (relative volumes of acid and base required for
neutralization) and record in the data table. The average titration volume will be used in the
calculation of the concentration of the acid.
Observations and Data:
Table 1: Titration Data
TRIAL 1
TRIAL 2
TRIAL 3
Initial base buret reading
Final base buret reading
Volume of sodium hydroxide used (VB)
Initial acid buret reading
Final acid buret reading
Volume of acetic acid used (VA)
Average volume of base required for titration (VB)
Molarity of NaOH
Average volume of acid used (VA)
Calculated Molarity of HC2H3O2
Calculations
1. Calculate the normality of the vinegar sample.
NA x VA = NB x VB
(Normality Acid x Volume Acid = Normality Base x Volume Base)
2. Calculate the molarity of the vinegar sample using the formula below
Molarity = Normality / total positive oxidation number
3. Calculate the grams of HC2H3O2 in the sample.
Molarity 
moles HC 2 H 3O 2
L of solution
1L  1000 mL
grams HC 2 H 3O 2  moles HC 2 H 3O 2  molar mass
4. Calculate the % acetic acid in the sample. The density of the vinegar is 1.0002 g/mL.
mass vinegar sample = 1.0002 g/mL x VA
% acid = (mass of HC2H3O2 / mass of vinegar solution) x 100%
5. Calculate your percent error for percent mass of acid.
Actual value = 5.0%
Discussion and Conclusion:
Follow guidelines outlined on the handout titled, “The Laboratory Record.”
Acid-Base Titration Analysis
Determination of the Acid Content in Vinegar
Table 1: Titration Data
TRIAL 1
TRIAL 2
Initial base buret reading
Final base buret reading
Total volume of sodium hydroxide used (VB)
Initial acid buret reading
Final acid buret reading
Total volume of acetic acid used (VA)
Average volume of base required for titration (VB)
Molarity of NaOH (stated in the materials list and in the procedures)
Average volume of acid used (VA)
Calculated Molarity of HC2H3O2
Calculations (based on average volume values from the trials listed above)
1. Calculate the normality of the vinegar sample.
2. Calculate the molarity of the vinegar sample using the formula below
3. Calculate the grams of HC2H3O2 in the sample.
TRIAL 3
4. Calculate the % acetic acid in the sample. The density of the vinegar is 1.0002 g/mL.
5. Calculate your percent error for percent mass of acid. (Actual value = 5.0%)
Discussion and Conclusion:
Final result: Concentration of acetic acid in vinegar is _________M and ________% acidity.
My percentage error in this lab was ________%
Table 2: Sources of error and suggestions for improvements
Detailed description of specific source of error
Description of a practical way of minimizing
(Note: NO “human error” or “calculation error”)
this error in a high school lab setting
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2
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