Laboratory Experiment pubs.acs.org/jchemeduc
Using Mole Ratios of Electrolytic Products of Water for Analysis of Household Vinegar: An Experiment for the Undergraduate Physical Chemistry Laboratory Rajeev B. Dabke* and Zewdu Gebeyehu Department of Chemistry, Columbus State University, Columbus, Georgia 31907, United States S Supporting Information *
ABSTRACT: A simple 3-h physical chemistry undergraduate experiment for the quantitative analysis of acetic acid in household vinegar is presented. The laboratory experiment combines titration concept with electrolysis and an application of the gas laws. A vinegar sample was placed in the cathode compartment of the electrolysis cell. Electrolysis of water generated OH−(aq) in the cathode compartment. Acetic acid in the vinegar sample was neutralized by electrolytically generated OH−(aq). Phenolphthalein was used as a visual indicator to detect the end point of the titration. The volume of O2(g) produced at the anode was monitored. The amount of OH−(aq) produced at the cathode was determined from the volume of O2(g) produced at the anode. The amount of acetic acid present in the sample was determined from the mole relationships between H+(aq), OH−(aq), and O2(g). The concentration of acetic acid in household vinegar was determined to be 4.8%. The experimentally determined percentage of acetic acid was in agreement with the manufacturer’s label and the volumetric titration results. The electrolysis cell was directly powered by a 9 V battery. Preparation of an electrolysis cell, experimental procedure, and results of the electrolytic titration are presented. KEYWORDS: First-Year Undergraduate/General, Laboratory Instruction, Physical Chemistry, Hands-On Learning/Manipulatives, Acids/Bases, Dyes/Pigments, Electrolytic/Galvanic Cells/Potentials, Gases, Oxidation/Reduction, Stoichiometry
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coulometric methods are composed of a constant current source and an ammeter.27 Coulometric titration experiments demand construction of a constant current supply28 or application of an expensive coulometric autotitrator.29 These stringent requirements may restrict the practice of coulometric titrations in an undergraduate laboratory. In this article, we present an easy-to-construct electrolysis titration cell in which coulometric measurement of electrical charge is eliminated. The cell is directly powered by a 9 V battery and does not involve any electronic circuit or an ammeter. The aim of the experiment is to determine the percentage of acetic acid in household vinegar. In this experiment, the volume of oxygen gas produced at the anode is monitored. The perfect gas equation is applied to determine the amount of oxygen gas produced at the anode. Faraday’s laws of electrolysis indicate that in an electrolysis process, the same quantity of electrical charge is responsible for producing the anode and cathode compartment products. It is evident from eqs 1 and 2that the amount of O2(g) produced at the anode and the amount of OH−(aq) produced at the cathode are in 1/2:2 or 1:4 mole ratio, respectively. This amount relationship is used to quantify the amount of OH−(aq) produced in the cathode compartment from the amount of O2(g) produced at the anode. The stoichiometric relation (eq
lectrolysis of water and stoichiometry of redox reactions are essential to high school and undergraduate chemistry courses. Many reports1−26 demonstrate the electrolysis of water and its pedagogical significance in chemistry. Traditionally, Hofmann apparatus or similar handmade devices have been used to demonstrate the relative proportion of gases produced as a result of the electrolysis process.1−3,6,8−10,12,13,17,21−23 Several indicators have been used to colorfully illustrate11,15,16,19,24−26 the pH changes during the electrolysis process. The electrolytic reactions are anode (oxidation): H 2O(l) → 2H+(aq) + 1 2 O2 (g) + 2e−
(1)
cathode (reduction): 2H 2O(l) + 2e− → 2OH−(aq) + H 2(g)
(2)
There are relatively few reports of electrolysis involving quantitative assessment of the gases produced at the respective electrodes.2,20−23 These quantitative experiments include determining Avogadro’s number,2,20 optimizing the electrolysis conditions,21,22 and determining the fundamental electronic charge.23 Silver-based coulometry played an important role in the early development of electrolytic chemical analysis.4,7,18 In these methods, the quantities of electrolytic products were determined exclusively from the mass of silver metal deposited on the electrode. However, these methods involved expensive and large-sized precious metal electrodes. Relatively new © 2012 American Chemical Society and Division of Chemical Education, Inc.
Published: July 13, 2012 1198
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3) is used to determine the percentage of acetic acid present in the household vinegar.
→ CH3COO (aq) + H 2O(l)
RESULTS
Electrolysis
CH3COOH(aq) + OH−(aq) −
Laboratory Experiment
The volume of O2(g) produced at the anode linearly changed with the volume of vinegar added to the cathode compartment (Figure 2). This linear dependence indicated the quantitative
(3)
The pedagogical objectives of this experiment are • to relate the perfect gas equation to quantitative analysis, • to apply amount relationships to quantify the active ingredients in commercial household vinegar sample, • to explore the anodic and cathodic reaction products and use them in a chemical analysis.
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EXPERIMENTAL OVERVIEW An electrolysis cell (Figure 1) was made of two polytetrafluoroethylene (PTFE) beakers. Platinum wires served as anode and
Figure 2. A plot of the volume of O2(g) produced at the anode versus volume of vinegar placed in the cathode compartment. The barometric pressure in the laboratory was 758.2 mm of Hg and the temperature of the electrolyte was 22.0 °C. The average height of the electrolyte column was 7.3 cm.
relation between the products generated at the cathode and anode compartments. The slope of the plot (Figure 2) was 5.1 (R2 = 0.99). The slope signified that 5.1 mL of oxygen gas was produced per unit volume of vinegar. From the perfect gas equation and from the stoichiometric relations presented in eqs 1 and 2, the amount of O2(g) and the amount of OH−(aq) produced in the cathode compartment were calculated. From the amount of OH−(aq), the percentage of acetic acid was determined to be 4.8%. The percentage of acetic acid determined from Figure 2 was in agreement with the manufacturer’s label (5%) and with the volumetric titration results (5.2%).
Figure 1. (A) Schematic diagram of the electrolysis cell: (a) PTFE beakers serving as cell compartments, (b) 0.2 M KNO3 as electrolyte, (c) platinum cathode, (d) hydrogen gas bubbles produced at the cathode, (e) Agar−KNO3 salt bridge, (f) Parafilm seal covering the tip, (g) graduated pipet, (h) oxygen gas bubbles produced at the anode, (i) platinum wire anode, (j) 9 V battery, (k) magnetic stirrer bar, and (l) switch. (B) Setup of the electrolysis cell. Cathode compartment shows pink coloration of phenolphthalein at the end point of the titration.
Optional Experiment
In a separate experiment, the pH of electrolyte containing a known volume of vinegar was monitored as a function of the volume of oxygen gas produced (Figure S1). A sharp increase in the pH determined the stoichiometric end point of the titration of acetic acid versus electrolytically produced OH−(aq). The percentage of vinegar determined from this plot was 4.8%. The quantitative results obtained from Figure 2 and Figure S1 were in agreement. This experiment was optional and the details are available in the Supporting Information.
cathode. The compartments were connected by an agar−KNO3 salt bridge. A known volume of vinegar was placed in the cathode compartment and the contents were magnetically stirred. A direct current was passed from a 9 V battery. Cathode compartment generated OH−(aq). Phenolphthalein was used as a visual indicator to determine the end point of the titration. An inverted glass pipet was used to measure the volume of oxygen gas produced at the anode. Aqueous potassium nitrate served as an electrolyte. A complete description of experimental procedures, instructions for the students, step-by-step calculations, and a post-laboratory exercise are included in the Supporting Information. Three undergraduate students performed the experiment in a 3-h physical chemistry laboratory and the data obtained by the students is reported in this article.
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DISCUSSION Additional experiments were run by the three undergraduate students to examine the experimental procedure. A blank trial ensured there was no measurable interference from the electrolyte and from dissolved carbon dioxide gas. In this trial, no vinegar was added to the electrolyte. The electrolyte instantly turned pink as a current supply from the battery was started. No measurable amount of O2(g) was produced in this trial. A blank trial was also performed after the vinegar titration to confirm there was no leakage of H+(aq) or OH−(aq) through the salt bridge during the titration. The current passing through the cell was in the 15−18 mA range.
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HAZARDS Safety glasses must be worn while performing the experiment. Potassium nitrate can cause eye and skin irritation. Labeled waste containers must be used to collect the waste chemicals. Because of the flammable nature of the H2(g), the titration should be performed in a fume hood. 1199
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Laboratory Experiment
Columbus State University’s STEM faculty development minigrant program.
The percent of acetic acid in the vinegar was determined from the volume of O2(g). Fluctuations in current were not relevant as the reagent concentration was not determined from the electrical charge. A typical tritration of 0.200 mL of vinegar took about 16 min. A 10 mL graduated pipet with 1/10 subdivisions was used for the volumetric measurement of oxygen. This pipet enabled the measurement of 0.1 mL or larger volumes of oxygen. The minimum measurable volume of oxygen determined the lower limit of detection of the analyte. Narrower and more sensitive pipets obstructed the flow of oxygen bubbles emerging from the electrode surface. Because no reagent was added from a buret during the titration, the preparation and standardization of reagents were completely excluded. There was no change in the volume of the electrolyte during the progress of the titration. The water vapor pressure added 3% to the O2(g) pressure. The descending level of electrolyte gradually dropped the hydrostatic pressure in the pipet. The average height of the column was used to determine the hydrostatic pressure. However, the hydrostatic pressure of the column contributed only 1% to O2(g) pressure. The pressure of O2(g) in equilibrium with the atmospheric pressure was corrected for both the water vapor pressure30 and the hydrostatic pressure.20 O2(g) was treated as a perfect gas for the amount calculation. No measurable change in the temperature of the cell contents during the titration was noticed. The average concentration of acetic acid determined from six trials was 4.8% with a standard deviation of ±8.9 × 10−2.
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CONCLUSION Undergraduate students with a prior knowledge of acid−base reactions can perform the electrolytic titration. The agreement between electrolytic titration results, manufacturer’s label, and volumetric titration results underlines the accuracy of the experiment. An easy-to-construct cell, elimination of electronic equipment, and study of a household product make this experiment relevant, cost-effective, and suitable for the undergraduate laboratory. Application of the perfect gas equation and quantitative mole relations directly addresses the pedagogical objectives. The pedagogy is applicable to the undergraduate quantitative analysis or physical chemistry laboratory course.
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REFERENCES
ASSOCIATED CONTENT
S Supporting Information *
Information for the instructors; handout for the students including a postlab exercise. This material is available via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS We thank James O. Schreck and Daniel Blumling for helpful discussions. We thank the reviewers of the manuscript for their comments and suggestions. We thank Claire Eunhye Cho, Heather Boyette, and Taralynn Williams for performing the experiment. We acknowledge the financial support from the 1200
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