Determining the Concentration of Acetic Acid

Titration for Acetic Acid in Vinegar

Hands-On Labs, Inc.
Version 42-0208-00-02

Lab Report Assistant

This document is not meant to be a substitute for a formal laboratory report. The Lab Report Assistant is simply a summary of the experiment’s questions, diagrams if needed, and data tables that should be addressed in a formal lab report. The intent is to facilitate students’ writing of lab reports by providing this information in an editable file which can be sent to an instructor.

Exercise 1: Determining the Concentration of Acetic Acid

Data Table 1. NaOH Titration Volume.

Initial NaOH Volume (mL)Final NaOH Volume (mL)Total volume of NaOH used (mL)
Trial 1927
Trial 290.58.5
Trial 3936
Average Volume of NaOH Used (mL) :

Data Table 2. Concentration of CH3COOH in Vinegar.

Average volume of NaOH used (mL)Concentration CH3COOH in vinegar (mol/L)% CH3COOH in vinegar
7.160.7166

Concentration CH3COOH in Vinegar = 0.00716L X 0.5/L X (1Mol CH3COOH/ 1 Mol NaOH) X 1/ 0.005L

Concentration of CH3COOH in Vinegar = 0.716Mol

Questions

A. The manufacturer of the vinegar used in the experiment stated that the vinegar contained 5.0% acetic acid. What is the percent error between your result and the manufacturer’s statement?

% Error = (6 – 5 / 5 ) x 100

= 20%

B. What challenges would you encounter with the titration if you had used apple cider vinegar or balsamic vinegar as the analyte instead of white vinegar? The white vinegar is made up of pure vinegar so the results will be more accurate than the apple cider vinegar or balsamic vinegar, the later has other “impurities” in them. The apple and balsamic has other properties which will interfere with the titration.

C. How would your results have differed if the tip of the titrator was not filled with NaOH before the initial volume reading was recorded? Explain your answer.

D. How would your results have differed if you had over-titrated (added drops of NaOH to the analyte beyond the stoichiometric equivalence point)?

E. If a 7.0 mL sample of vinegar was titrated to the stoichiometric equivalence point with 7.5 mL of 1.5M NaOH, what is the mass percent of CH3COOH in the vinegar sample?

F. Why is it important to do multiple trials of a titration, instead of only one trial?

Titration for Acetic Acid in
Vinegar
Hands-On Labs, Inc.
Version 42-0208-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in which
to complete the exercise.

Experiment Summary:

You will learn about acetic acid (CH3COOH) and how to
determine the concentration of acetic acid in vinegar
through titration. You will explore the concepts of
stoichiometric equilibrium, concentration, molarity,
indicators, and mass percent.

EXPERIMENT

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Learning Objectives
Upon completion of this laboratory, you will be able to:

● Define titration, titrant, analyte, and equivalence (or stoichiometric) point.

● Discuss why phenolphthalein is an effective pH indictor, and how it works.

● Describe the process and purpose of titration.

● Write a balanced chemical equation representing the stoichiometric equivalence point of a
reaction.

● Describe how strong and weak acids differ.

● Apply titration techniques to investigate acetic acid in commercial vinegar.

● Determine the molar concentration of acetic acid in commercial vinegar.

● Calculate the average concentration and the percent (%) concentration of acetic acid in vinegar.

Time Allocation: 2.5 hours

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Experiment Titration for Acetic Acid in Vinegar

Materials
Student Supplied Materials

Quantity Item Description
1 Bottle of distilled water
1 Dish soap
1 Pair of scissors
1 Roll of paper towels
1 Sheet of white paper
1 Source of tap water

2-6 Textbooks

HOL Supplied Materials

Quantity Item Description
1 Glass beaker, 100 mL
1 Graduated cylinder, 25 mL
1 Pair of safety goggles
1 Pair of gloves
1 Stopcock
1 Syringe, 10 mL
1 Test tube clamp
1 Test tube cleaning brush
1 Experiment Bag: Titration for Acetic Acid in Vinegar

1- Phenolphthalein solution, 1%, 0.5 mL
1- Vinegar, 20 mL in dropper bottle
1- Pipet, short stem
1- Sodium hydroxide, 0.5 M, 30 mL

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Titration for Acetic Acid in Vinegar

Background
Acetic Acid

Have you ever opened a bottle of vinegar and noticed the intense scent? Now think about the
sour taste you get while nibbling on a bag of salt and vinegar potato chips. You might be surprised
to learn that both the intense scent of vinegar and sour taste of salt and vinegar chips are the
result of acetic acid, the chemical component of vinegar. Acids are molecular substances that
ionize in water to release H+ ions. Acetic acid (CH3COOH) is a weak acid, which means that only
a small percentage of acetic acid molecules ionize when dissolved in water (only the hydrogen
atom bonded to the oxygen ionizes). In contrast, the acidic hydrogen atoms in a strong acid are
completely ionized in water. All acids, weak or strong, form a conjugate base and some number
of protons when dissolved in water. See Figure 1.

Figure 1. Strong and weak acids. A. The weak acid, acetic acid (CH3COOH), only partially ionizes
in water. Notice that of the four total hydrogen molecules in acetic acid, only one H+ ion is

produced in solution. B. The strong acid, nitric acid (HNO3), completely ionizes in water.

The amount of acetic acid in common table vinegars, apple cider vinegar, balsamic vinegar, white
vinegar, and sherry vinegar, varies from approximately 4-6%. However, some varieties of vinegar,
such as pickling vinegar or distilled vinegar, contain approximately 5-8% acetic acid.

Titration

The exact concentration of acetic acid in a bottle of vinegar can be determined through titration
with a strong base. Titration is a quantitative, volumetric technique where a solution of a known
concentration (titrant) is added to a solution of an unknown concentration (analyte) until the
equivalence point is reached. See Figure 2. The equivalence point of a titration, also known as a
stoichiometric point or end point, is the moment in a titration where exactly enough titrant has
been added to completely react with the analyte. In an acid-base titration, the equivalence point
can be identified through the use of a pH indicator.

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Experiment Titration for Acetic Acid in Vinegar

Figure 2. Laboratory titration apparatus. © Alexander Raths

A pH indicator is a substance that changes color when the pH of a solution changes, allowing
scientists to qualitatively measure the moment when the analyte has completely reacted
with the titrant. A common indicator for a titration between a weak acid and a strong base is
phenolphthalein. See Figure 3. Phenolphthalein is a pH indicator, which turns bright-pink at a basic
pH of 8.2 and higher, allowing for equivalence points in titrations to be marked by the analyte
changing in color from colorless to bright pink.

Figure 3. Phenolphthalein added to solution. © Kesu

When designing a titration between an acid and a base, it is important to know the reaction
between titrant and analyte (chemical equation), the exact stoichiometric equivalence point
(balanced chemical equation), and the concentration of the titrant. To determine the exact
concentration of acetic acid in vinegar, a titration between acetic acid (CH3COOH) and sodium
hydroxide (NaOH) will be performed, using phenolphthalein as the indicator. The balanced chemical
equation for the reaction between CH3COOH and NaOH is shown below:

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Experiment Titration for Acetic Acid in Vinegar

From the balanced equation, we know that the reaction between acetic acid and sodium hydroxide
is 1:1, thus the number of moles of NaOH required to cause a change in the indicator color is equal
to the number of moles of CH3COOH in the vinegar. The stoichiometric equivalence point (end
point) of the titration will be visualized when the phenolphthalein causes the solution to change
from a colorless solution to a bright-pink colored solution.

The titration will provide the number of moles of CH3COOH present in the vinegar sample. In order
to determine the %CH3COOH in the sample, the mass percent (grams of acetic acid/grams of
vinegar) will be determined using the molar mass of CH3COOH (60.05g CH3COOH/1 mol CH3COOH)
and the density of vinegar (1.00 g/mL).

Titration is used in many
different ways to measure chemical

concentrations in solutions: Nutritional analysis
often uses titration methods to measure acidity in

a food (for example, orange juice); Blood testing and
pregnancy tests may use methods of titration to measure

chemical levels; Aquarium water testing uses titration
methodologies to ensure a healthy ecosystem for aquatic
pets; Winemakers may use fairly simple titration kits to

measure the acidity of wine; The development of medications
(pharmacology) by pharmaceutical companies frequently

involves the use of specialized titration equipment to
accurately measure chemical quantities; Acid rain evaluation

employs titration processes to quantify the degree of
contamination in natural rain water or snow;

Wastewater analysis often uses specialized titration
equipment to analyze contamination, and thus
determine the requirements for filtering and

cleaning.

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Experiment Titration for Acetic Acid in Vinegar

Exercise 1: Determining the Concentration of Acetic
Acid
In this exercise, you will determine the concentration of acetic acid (CH3COOH) in the vinegar
provided in your lab kit.

Note: Please read all steps and safety information before starting the procedure.

Procedure

1. Gather the test tube holder, small stopcock, 10-mL syringe (titrator), and either 2 thick textbooks
and the lab kit box or 5-6 thick textbooks. See Figure 4.

Figure 4. Titrator and small stopcock.

2. Remove the plunger from the titrator and place it back in your lab kit box.

3. Attach the stopcock to the tip of the titrator by placing the larger, clear, plastic end of the
stopcock into the tip of the titrator and then twist the stopcock into place. The stopcock should
fit tightly into the titrator so that the liquid will not leak. See Figure 5.

Figure 5. Fitting the stopcock into the titrator.

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Experiment Titration for Acetic Acid in Vinegar

4. Stack the 5 textbooks or stack 2 textbooks on top of the lab kit box.

5. Clamp the test tube holder around the middle of the titrator and slide the long end under the
top 2 books in the stack. Place a sheet of white paper next to the bottom of the stack and set
the 100 mL beaker on the sheet of white paper. The end of the stopcock should be located
near the top of the beaker, approximately 1 cm above to 1 cm below the top of the beaker.
See Figure 6.

Figure 6. Titration setup.

Note: It is important that the placement of the titrator allows for the white knob to be easily adjusted.
If this is not the case, then either adjust the location of the books in the stack or slightly adjust where
in the test tube clamp the titrator is located.

6. Use the pipet to fill the titrator with 7 – 9 mL of distilled water.

Note: It is important that you use distilled water for this step and not tap water.

7. Using both hands, one hand on the titrator and your other hand on the stopcock, practice
releasing water from the titrator into the beaker. The goal is to be comfortable releasing only
1 drop at a time from the titrator. See Figure 7.

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Experiment Titration for Acetic Acid in Vinegar

Figure 7. Proper hand positioning for titration. When the open circle is facing you, the titrator is
closed, when the open circle is directly under the titrator spout, the titrator is open and liquid

will flow.

8. When you are comfortable using the titrator, pour the water in the beaker down the drain,
remove the titrator from the test tube clamp, and remove the stopcock from the titrator.
Thoroughly dry each of these 3 items with paper towels.

9. When all items are completely dry, reassemble the titration setup, as shown in Figure 6.

10. Put on your safety gloves and goggles.

11. With the stopcock in the closed position, fill the titrator with 9 – 10 mL of the 0.5M NaOH.

12. Move the beaker away from the titrator and place a crumpled paper towel directly below the
titrator.

13. Using the stopcock, allow a few drops of the NaOH to flow through the titrator into the paper
towel. This will fill the tip of the titrator with NaOH solution and remove any air bubbles from
the titrator.

14. Place the paper towel with the NaOH drops into the trash and reposition the clean, dry 100
mL beaker back in the titration setup, under the titrator.

15. Use the graduated cylinder to measure exactly 5 mL of vinegar.

16. Pour the 5 mL of vinegar into the completely dry 100 mL beaker.

17. Cut off the tip of the pipet with scissors and add 2 drops of phenolphthalein to the 5 mL of
vinegar in the beaker.

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Experiment Titration for Acetic Acid in Vinegar

18. Carefully swirl the mixture in the beaker to ensure that the indicator is incorporated into the
vinegar; the solution will be colorless and clear.

19. Read the volume of NaOH in the titrator and record in Data Table 1 of your Lab Report Assistant
under “Initial NaOH Volume (mL),” next to “Trial 1”.

20. Open the stopcock and add 1 drop of NaOH to the colorless and clear vinegar sample in the
beaker. After the drop is added, gently swirl the beaker and observe the color for 5 seconds.

Note: It is important to add the NaOH 1 drop at a time to avoid overshooting the titration.

21. Continue adding NaOH to the beaker, 1 drop at a time, swirling and observing after each drop
until the color changes to a bright-pink color for at least 5 seconds. See Figure 8.

Figure 8. Endpoint of titration. The titration is complete when the color changes to a bright pink
for at least 5 seconds.

22. Read the volume of the NaOH solution remaining in the titrator and record this volume in Data
Table 1 under “Final NaOH Volume (mL)”, next to “Trial 1”.

23. Determine the total volume of NaOH used by subtracting the final NaOH volume from the
initial NaOH volume and record the total volume in Data Table 1.

24. Leave the titrator assembly intact. You will need it for future titrations in this experiment.

25. Pour the contents of the beaker down the drain and flush the drain with water. Thoroughly
wash the beaker with soap and water to remove all of the vinegar/NaOH/indicator solution
from the beaker. When the beaker is clean, rinse the beaker with distilled water and then
thoroughly dry.

26. Place the clean and thoroughly dry beaker back in the titration setup, under the titrator.

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Experiment Titration for Acetic Acid in Vinegar

27. If necessary, add more NaOH to the titrator.

Note: It is only necessary to add more NaOH to the titrator if there is less than 1 mL more than the
total volume of NaOH used in the previous trial. For example, if the total volume of NaOH used in
Trial 1 was 2.1 mL, then there needs to be at least 3.1 mL of NaOH in the titrator.

28. Repeat Steps 15 through 27 two additional times (Trial 2 and Trial 3), using the vinegar provided
in the lab kit.

29. Average the results from the three trials and record in Data Table 1 and Data Table 2 of your
Lab Report Assistant.

30. Using the following equation, determine the average concentration (moles per liter) of acetic
acid (CH3COOH) present in your vinegar. Record the concentration in Data Table 2.

31. Using the following equation, determine % concentration (mass percent) of acetic acid
(CH3COOH) in the vinegar and record it in Data Table 2.

Note: The density of vinegar is 1.00 g/ml and the molecular mass of vinegar is 60.05 g/mol.

Cleanup:

32. Properly dispose of remaining chemicals.

33. Wash and dry all equipment and return to the lab kit.

Questions
A. The manufacturer of the vinegar used in the experiment stated that the vinegar contained 5.0%

acetic acid. What is the percent error between your result and the manufacturer’s statement?

B. What challenges would you encounter with the titration if you had used apple cider vinegar
or balsamic vinegar as the analyte instead of white vinegar?

C. How would your results have differed if the tip of the titrator was not filled with NaOH before
the initial volume reading was recorded? Explain your answer.

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Experiment Titration for Acetic Acid in Vinegar

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Experiment Titration for Acetic Acid in Vinegar

D. How would your results have differed if you had over-titrated (added drops of NaOH to the
analyte beyond the stoichiometric equivalence point)?

E. If a 7.0 mL sample of vinegar was titrated to the stoichiometric equivalence point with 7.5 mL
of 1.5M NaOH, what is the mass percent of CH3COOH in the vinegar sample?

F. Why is it important to do multiple trials of a titration, instead of only one trial?

GENERAL CHEMISTRY 1

GENERAL CHEMISTRY 3

Course title: General Chemistry II

Course Code: CHM 151

Student Name:

Submission Date: th April 2017

8.Write the ionic equation for dissolution and the solubility product (Ksp) expression for each of the following slightly soluble ionic compounds: (a) PbCl2 (b) Ag2S (c) Sr3(PO4)2 (d) SrSO4

14.Assuming that no equilibria other than dissolution are involved, calculate the molar solubility of each of the following from its solubility product:

(a) Ag2SO4

(b) PbBr2

(c) AgI

(d) CaC2O4∙H2O

28. The following concentrations are found in mixtures of ions in equilibrium with slightly soluble solids. From the concentrations given, calculate Ksp for each of the slightly soluble solids indicated:

(a) AgBr: [Ag+] = 5.7 × 10–7 M, [Br–] = 5.7 × 10–7 M

(b) CaCO3: [Ca2+] = 5.3 × 10–3 M, [CO3 2−] = 9.0 × 10–7 M

(c) PbF2: [Pb2+] = 2.1 × 10–3 M, [F–] = 4.2 × 10–3 M

(d) Ag2CrO4: [Ag+] = 5.3 × 10–5 M, 3.2 × 10–3 M

(e) InF3: [In3+] = 2.3 × 10–3 M, [F–] = 7.0 × 10–3 M

54. Calculate the molar solubility of AgBr in 0.035 M NaBr (Ksp = 5 × 10–13).

CHAPTER 16

4.A helium-filled balloon spontaneously deflates overnight as He atoms diffuse through the wall of the balloon. Describe the redistribution of matter and/or energy that accompanies this process.

12.Arrange the following sets of systems in order of increasing entropy. Assume one mole of each substance and the same temperature for each member of a set.

(a) H2(g), HBrO4(g), HBr(g)

(b) H2O(l), H2O(g), H2O(s)

(c) He(g), Cl2(g), P4(g)

16.Predict the sign of the entropy change for the following processes.

(a) An ice cube is warmed to near its melting point.

(b) Exhaled breath forms fog on a cold morning.

(c) Snow melts

16.Predict the sign of the entropy change for the following processes.

(a) An ice cube is warmed to near its melting point.

(b) Exhaled breath forms fog on a cold morning.

(c) Snow melts

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