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E XPERIMENT A7: VINEGAR T ITRATION
Learning Outcomes
Upon completion of this lab, the student will be able to:
- Prepare a solution of primary standard
- Determine the molar concentration of a solution of an acid or base using data obtained from titration
- Measure the amount of acetic acid in a solution of vinegar
Introduction
The molar concentration of an acid or a base can be determined by the method of titration. In a titration, a solution of known concentration is slowly added to a known volume of a solution of unknown concentration until the two have completely reacted with each other. There are many different kinds of titration; the type used in this experiment is called an acid-base titration.
Consider for instance, the reaction between aqueous solutions of sulfuric acid and potassium hydroxide. The balanced chemical equation for the reaction is shown below:
H 2 SO4(aq) + 2KOH (aq) K 2 SO4(aq) + 2H 2 O(l)
Assume that the molar concentration of KOH is 0.100 M and that of H 2 SO 4 is unknown. In order to determine the molar concentration of H 2 SO 4 , one would need to add the KOH solution to a known volume of H 2 SO 4.
In a certain experiment, the KOH was added to 10.00 mL of H 2 SO 4. The volume of KOH needed to completely react with the H 2 SO 4 will enable the determination of the molarity of H 2 SO 4. In the same experiment, assume that 13.75 mL of KOH was needed to completely react with the H 2 SO 4. The molarity of the acid is calculated as follows:
Molarity of H 2 SO 4 =
As seen from the above calculation, the stoichiometric ratio between the two reactants is the key to the determination of the molarity of the unknown solution.
In order to conduct the above experiment, typically the H 2 SO 4 is in an Erlenmeyer flask, and the KOH is in a burette. The KOH is added one drop at a time from the
burette into the acid solution with constant stirring to ensure that the reagents combine and react.
Vinegar Titration
In this experiment, a solution of vinegar has been provided for analysis. The active ingredient in vinegar is acetic acid, CH 3 COOH. In order to determine the amount of acetic acid in the vinegar, the acetic acid will be titrated with a solution of known concentration of sodium hydroxide.
The chemical reaction between acetic acid and sodium hydroxide is given below:
CH 3 COOH (^) (aq) + NaOH (^) (aq) CH 3 COONa (^) (aq) + H 2 O(l)
The balanced chemical equation shows that one mole of CH 3 COOH reacts with exactly one mole of NaOH. The experiment is performed by adding NaOH of known molarity to a known volume of vinegar until the reaction is complete.
Determining the completion of an acid/base reaction
In order to obtain the molarity or moles of the unknown reactant the solution whose concentration is known (in this experiment that would be NaOH), must be added until the reaction is complete. This means, exactly one mole of NaOH must be added to one mole of acetic acid. This point in the titration is called the Equivalence Point.
The equivalence point is defined as that point in the titration when stoichiometrically equal amounts of acid and base are present. In the CH 3 COOH/NaOH titration, that would be when one mole of NaOH has been added to one mole of CH 3 COOH. In the H 2 SO 4 /KOH example shown previously, that would be when two moles of KOH have been added to one mole of H 2 SO 4. Therefore the equivalence point depends on the reaction stoichiometry.
At the beginning of the titration, the solution in the Erlenmeyer flask is acidic. As the base is added, it completely reacts with the acid and the solution in the Erlenmeyer flask continues to be acidic. But, at the equivalence point, the acid has completely reacted with the base. If even one tiny drop of base is added beyond that needed to arrive at the equivalence point, the solution in the Erlenmeyer flask is basic. This difference in the acid/base property of the solution in the Erlenmeyer flask is used to visually determine the end of the titration.
An indicator is a chemical substance whose color depends on the acid/base property of the medium it is present in. Phenolphthalein is an indicator, which is colorless in an acidic medium and has a pink color in a basic medium.
In this titration, a few drops of phenolphthalein should be added to the acid in the Erlenmeyer flask. The solution will remain colorless until the equivalence point.
Preparation of primary standard solution
The moles of acetic acid in vinegar will be obtained by titrating the acetic acid with a solution of sodium hydroxide of known molarity. The molarity of the sodium hydroxide solution will be determined by titrating it with a solution of KHP of known molarity. Therefore it should be apparent that the accuracy of the results of the experiment depends on the accuracy of the molarity of the KHP. The KHP solution should be prepared using a volumetric flask. When using a volumetric flask, care must be taken to avoid crossing the calibrated mark.
Experimental Design
A solution of sodium hydroxide whose concentration is known approximately will be provided for this experiment. The exact molarity of this sodium hydroxide solution will be determined by titrating it with a solution of KHP that must be prepared by the experimenter. Once the molarity of the sodium hydroxide solution is determined, it will then be used to titrate the acetic acid in the vinegar. Phenolphthalein is used as an indicator in both of these titrations.
Reagents and Supplies
1.0 M aqueous sodium hydroxide, solid potassium hydrogen phthalate, commercial vinegar solution, 1% and 0.25% phenolphthalein solution
25-mL volumetric flask
(See posted Material Safety Data Sheets)
- Dispense a little over 1-mL of KHP solution into a small Erlenmeyer flask and record the exact final burette reading.
- Rinse the sides of the Erlenmeyer flask with deionized water from a squirt bottle.
- Titrate the KHP solution in the Erlenmeyer flask with the NaOH solution until a permanent pale pink color is obtained. After the addition of each drop of NaOH, be sure to swirl the Erlenmeyer flask thoroughly to ensure mixing of the reagents. In case any NaOH solution falls on the side of the Erlenmeyer flask, rinse the sides of the flask with deionized water.
- Repeat steps 10-13 at least two more times.
P ART 2: TITRATION OF THE ACETIC ACID IN VINEGAR WITH THE STANDARDIZED SODIUM HYDROXIDE SOLUTION
- Use the same NaOH solution as in part 1 of this experiment.
- Obtain approximately 10-mL of commercial vinegar solution.
- Clamp two microburettes to a burette stand and label one burette as “NaOH” and the other burette as “vinegar”.
- Condition each burette with the respective reagent.
- Record the initial burette readings of both the burettes.
- Obtain a clean, dry, small Erlenmeyer flask.
- Dispense approximately 0.1-mL of vinegar into the Erlenmeyer flask.
- Record the mass of the Erlenmeyer flask with the vinegar.
- Add one drop of 0.25%-phenolphthalein solution into the Erlenmeyer flask containing the vinegar.
- Rinse the sides of the Erlenmeyer flask with deionized water.
- Titrate the contents of the Erlenmeyer flask with the NaOH solution until a permanent pale pink color is obtained (Refer to step 13 of part 1).
- Record the final burette reading of the NaOH burette.
- Repeat steps 5-13 two or three more times.
- Dispose all waste into appropriate waste disposal containers as instructed by your instructor.
P ART 2: TITRATION OF THE ACETIC ACID IN VINEGAR WITH THE STANDARDIZED SODIUM HYDROXIDE
SOLUTION
Vinegar
Trial 1 Trial 2 Trial 3 Trial 4
Initial burette reading (mL) Final burette reading (mL) Volume Used (mL)
NaOH solution
Trial 1 Trial 2 Trial 3 Trial 4
Initial burette reading (mL) Final burette reading (mL) Volume Used (mL)
Calculations
P ART 1: S TANDARDIZATION OF AQUEOUS SODIUM HYDROXIDE SOLUTION
Molar mass of KHP (C 8 H 5 O 4 K) =
Mass of KHP =
Moles of KHP =
Volume of KHP solution prepared = 25.00 mL = 0.02500 L
Molarity of KHP = =
Trial 1 Trial 2 Trial 3 Trial 4 Volume of KHP (liters) Volume of NaOH (liters)
Molarity of NaOH (show calculation for each trial): C 8 H 5 O 4 K + NaOH C 8 H 4 O 4 KNa + H 2 O
Trial 1
Trial 2
Trial 3
Trial 4
Molarity of NaOH
Trial 1
Trial 2
Trial 3
Trial 4
Average
Average mass percent of acetic acid in vinegar = _______________________
