Analisys of Vinegar via Titration

Note: This is not an experiment for CHM152 students – this is an example to illustrate the proper format for a lab report. • Pages 1-2 provide the experiment that was performed • Pages 3-4 show the sample lab report that a student would write and submit to be graded. *Formal Lab reports must be neatly handwritten or typed, or a combination of the two. Reports cannot be pages torn from your lab notebook. ANALYSIS OF VINEGAR VIA TITRATION INTRODUCTION In an acid-base neutralization reaction, an acid reacts with a base to produce water and a salt: HX (aq) + MOH (aq) acid ⇒ base H2O (l) + MX (aq) (1) salt The protons (H+) from the acid react with the hydroxide ions (OH–) from the base to form the water. The salt forms from the cation from the base and the anion from the acid. Because water is always formed, acids will always react with bases; whether the salt is soluble or insoluble does not determine whether the reaction occurs. In this experiment, you will use a dilute solution of sodium hydroxide, NaOH (aq) to determine the molar concentration of acetic acid present in a sample of vinegar. You will measure out a small volume of vinegar and use a buret to determine the volume of sodium hydroxide required to completely neutralize the vinegar. The process of slowly adding one solution to another until the reaction between the two is complete is called a titration. The reaction between acetic acid, HC2H3O2 (aq), and sodium hydroxide, NaOH (aq), is shown below: HC2H3O2 (aq) + NaOH (aq) ⇒ H2O (l) + NaC2H3O2 (aq) (2) When carrying out an acid-base neutralization reaction in the laboratory, you observe that most acid solutions and base solutions are colorless, and the resulting water and soluble salt solutions are also colorless. Thus, it is impossible to determine when a reaction has occurred, let alone when it is complete. To monitor the progress of a neutralization reaction, you use an acid-base indicator, a solution that changes color depending on the pH (or acid-content) of the solution. One commonly used indicator is phenolphthalein, which is colorless in acidic and neutral solutions and pink in basic (or alkaline) solution. During a titration, the indicator is added to the solution being titrated (the analyte). The titrant (or standard) is slowly added to the analyte until the endpoint, when the indicator changes color, signaling that the reaction between the two is complete. Note that phenolphthalein turns pink only when excess sodium hydroxide, NaOH (aq), has been added. The equivalence point in a titration is when there are equal numbers of moles of acid and base present in the solution. Thus, for the reaction between KHP and sodium hydroxide, at the equivalence point 1 of 4 moles of HC2H3O2 = moles of NaOH. (4) You should recognize that the end point is determined by observation of a color change in the lab while the equivalence point is a theoretical point one can calculate or estimate given the amount of acid or base present to titrate. PROCEDURE Analysis of Vinegar 1. Obtain ~50 mL of vinegar in your 100-mL beaker. Record the brand of vinegar and the percent acetic acid indicated on the label. 2. Condition a 10-mL volumetric pipet with the vinegar solution. If you are unfamiliar with the operation of a pipet, ask your instructor for help. Obtain and label three Erlenmeyer flasks #1, #2, and #3. Pipet 10 mL of vinegar to each of three Erlenmeyer flasks. Record the volume for each trial as 10.00 mL. Add ~25 mL of distilled water to each flask 3. Add 2-3 drops of phenolphthalein indicator to each of your vinegar solutions. 4. Obtain about 150 mL of the dilute NaOH solution, and record the molar concentration for the NaOH. Condition a 50-mL buret with the NaOH solution. 5. Fill the buret close to 0.00 mL with the dilute NaOH solution, and record the initial volume for trial #1 for vinegar on your Data Sheet. Slowly lower the stir bar into flask #1, and carry out the titration as before. Record the final volume of NaOH on your Data Sheet. 6. Repeat the titration for trials #2 and #3. 7. Dispose of all the solutions in the proper waste containers. Wash and rinse your buret and flasks, and return them to their proper places. BE SURE TO WASH AND DRY YOUR LAB BENCH AFTER COMPLETING THE EXPERIMENT TO REMOVE ALL TRACES OF ANY SPILLED CHEMICALS. CALCULATIONS 1. Calculate the molarity of the acetic acid in each vinegar solutions. Next, calculate the average molarity of the acetic acid for all three solutions. 2. Determine the density of the vinegar solution. 3. Calculate the mass percentage of acetic acid in the vinegar. QUESTIONS 1. Compare the mass percentage you determined with the mass percentage reported on the bottle of vinegar. PAGE 2 Title: Analysis of Vinegar via Titration Date: 08/15/2007 Name: Joe Student Partner: Sally Partner Purpose: The objective of this lab is to determine the molar concentration and mass percentage of acetic acid, HC2H3O2(aq), in a vinegar solution via titration using a standard solution of sodium hydroxide, NaOH(aq). Procedure: (Only if required) Obtain a 50-mL sample of vinegar. Record the brand and mass percentage of acetic acid reported on the vinegar bottle. Pipet 10-mL of vinegar into each of three Erlenmeyer flasks, and add a few drops of phenolphthalein indicator to each flask. Obtain about 150 mL of a dilute standard sodium hydroxide solution, NaOH(aq), and record the molar concentration. Clean and fill a buret with the NaOH solution. Titrate the three vinegar solutions to a light pink endpoint, recording the initial and final buret volumes required to titrate each sample. From the data you obtain in the three titrations, calculate the molar concentration of HC2H3O2 in each vinegar sample and the average molarity of HC2H3O2. Record the mass of an empty graduated cylinder. Add 10 ml of vinegar to the flask and record the mass. Determine the mass and the density of the vinegar solution. Use the average molarity of HC2H3O2 and the density of the vinegar solution to calculate the mass percent of HC2H3O2 in the vinegar. Data: Molarity of standard NaOH(aq): 0.2160 M Vinegar brand: Smiths Reported mass percent of HC2H3O2 in vinegar: 5.0% Table 1: Vinegar Titration: Final Volume of NaOH Initial Volume of NaOH Volume of NaOH used 10.00 ml vinegar used for each titration Trial 1 Trial 2 Trial 3 39.00 mL 38.95 mL 38.90 mL 0.00 mL 0.15 mL 0.10 mL 39.00 mL 38.80 mL 38.80 mL Table 2: Density of Vinegar Solution: Mass of Graduated Cylinder + vinegar 25.1753 g Mass of Empty 10-mL Graduated Cylinder 15.1256 g Mass of vinegar 10.0503 g Volume of vinegar used 10.00 mL 3 of 4 Sample Calculations: Molarity of HC2H3O2 in vinegar: Trial 1: 0.03900 L NaOH × 0.216 mol NaOH 1 mol HC2H3O2 1 × × = 0.842 M HC2H3O2 L 1 mol NaOH 0.01000 L Trials 2 and 3: see results table Average [HC2H3O2] = Density of vinegar: 0.842 M + 0.838 M + 0.838 M = 0.839 M 3 density = 10.0503 g mass = = 1.005 g/mL volume 10.00 mL Mass percent of HC2H3O2 in vinegar: 0.839 mol HC H O 60.052 g HC H O 2 2 3× 2 2 3 × mL solution × 100% = 5.01% 1000 mL solution mol HC H O 1.005 g solution 2 2 3 Vinegar Titration Results Molarity of HC2H3O2 in vinegar Average Molarity of HC2H3O2 Density of vinegar solution Mass percent of HC2H3O2 in vinegar Trial 1 Trial 2 Trial 3 0.842 M 0.838 M 0.838 M 0.839 M 1.005 g/mL 5.01% Questions: #1. The mass percentage of the acetic acid in the vinegar solution was found to be 5.01%, which is in close agreement with the manufacturer’s reported acetic acid content of 5.0%. Conclusions: In this experiment, a sample of Smiths brand vinegar was analyzed. Three trials were carried out, and the average molarity of the vinegar was determined to be 0.839 M. The mass percentage of the acetic acid, HC2H3O2(aq), in the vinegar was found to be 5.01%, which is only slightly higher than the manufacturer’s reported acetic acid content of 5.0%. One possible source of error was that NaOH was added past the endpoint for trial #1, for which 0.20 mL more NaOH was used compared to trials #2 and #3 even though the same volume of vinegar was titrated for all three trials. PAGE 4