Volumetric Titrations Using Electrolytically Generated Reagents for the

Apr 18, 2014 - Proposed electrolysis cell was directly powered by a 9 V battery eliminating the requirement of a coulometer. Students use this method ...
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Volumetric Titrations Using Electrolytically Generated Reagents for the Determination of Ascorbic Acid and Iron in Dietary Supplement Tablets: An Undergraduate Laboratory Experiment Christopher Scanlon, Zewdu Gebeyehu, Kameron Griffin, and Rajeev B. Dabke* Department of Chemistry, Columbus State University, Columbus, Georgia 31907, United States S Supporting Information *

ABSTRACT: An undergraduate laboratory experiment for the volumetric quantitative analysis of ascorbic acid and iron in dietary supplement tablets is presented. Powdered samples of the dietary supplement tablets were volumetrically titrated against electrolytically generated reagents, and the mass of dietary reagent in the tablet was determined from the volume of potassium hydrogen phthalate solution used as a primary standard. Proposed electrolysis cell was directly powered by a 9 V battery eliminating the requirement of a coulometer. Students use this method to quantify the active ingredient in dietary supplement tablets: iron and ascorbic acid (vitamin C). The experiment is suitable for the undergraduate quantitative analysis laboratory.

KEYWORDS: Second-Year Undergraduate, Upper-Division Undergraduate, Analytical Chemistry, Laboratory Instructions, Hands-On Learning/Manipulatives, Acids/Bases, Aqueous Solution Chemistry, Electrochemistry, Electrolytic/Galvanic Cells/Potentials, Reactions

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circuitry. The cell (Figure 1) is directly powered by a 9 V battery and does not involve any constant current electronic circuit or a milli-ammeter. The cell is housed in a test tube and relatively small volumes of electrolytes are required. Students with a prior knowledge of acid−base and redox titrations can perform these experiments. Thus, the two titration experiments presented in this paper can be implemented as quantitative analysis laboratory experiments. Quantifying the household products and connecting the pedagogy of electrolytic and chemical reactions make this experiment suitable and interesting to students in an undergraduate analytical laboratory. Pedagogical objectives of this experiment are · to use the electrolytic products for volumetric quantitative analysis, · to apply the stoichiometric relations in chemical equations, and · to explore the practical applications of redox reactions.

onventional volumetric titrations involving I2(aq) and Ce4+(aq) as reagents necessitate independent standardizations and require large volumes of reagents.1,2 In a coulometric titration, a reagent is produced in an electrolytic oxidation or reduction process. The species to be determined is titrated against the electrolytically produced reagent in a secondary chemical reaction. Faraday’s laws of electrolysis are then applied to quantify the reagent and the analyte species from the charge associated with the electrolysis process.3 Because the amount of reagent produced in a coulometric titration is directly determined from the charge passing through the cell, no standardization of reagents is required. Although less waste chemicals are produced in the coulometric methods, maintaining a constant current or monitoring the charge passing through an electrolysis cell is critically important.3−7 This stringent requirement may restrict the practice of coulometric titrations in an undergraduate quantitative analysis laboratory. The experiment described here couples two titrations of electrolytically produced reagents avoiding the need of a coulometer. This electrolytic titration method demonstrates the fact that in completing the circuit, each electron may take part in more than one reaction and by monitoring the progress of one reaction, the progress of the other can be followed. Students use this method to quantify the active ingredients in two dietary supplement tablets: iron and ascorbic acid (vitamin C). The electrolysis cell presented in this paper offers the advantages of coulometric methods, as well as simplifying the © 2014 American Chemical Society and Division of Chemical Education, Inc.

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EXPERIMENTAL OVERVIEW A unique electrolytic titration method using the primary standard substance 8 potassium hydrogen phthalate (KHC8H4O4(aq) or KHP(aq)) and electrolytically produced reagents is presented in this paper. The reagents produced in the cathodic and anodic electrolytic reactions and their Published: April 18, 2014 898

dx.doi.org/10.1021/ed400630u | J. Chem. Educ. 2014, 91, 898−901

Journal of Chemical Education

Laboratory Experiment

Table 1. Summary of Electrolytic and Chemical Reactions in the Electrolysis Cell

V battery is passed, electrolytic oxidation of KI(aq) produces I2(aq). Electrolytically produced I2(aq) (in form of I3−(aq)) reagent chemically oxidizes ascorbic acid in the sample (Table 1, eqs 3 and 4) and the end point of this titration is determined by a visual color change of the starch indicator from colorless to black (Figures S1A,B). The electrolysis is promptly stopped at the end point. Concurrent electrolysis of water in the cathode compartment produces OH−(aq) (Table 1, eq 1). Production of OH−(aq) in the presence of phenolphthalein indicator causes pink coloration of the electrolyte (Figure S1B). The electrolytically produced OH−(aq) is titrated against KHP(aq) to determine the amount of OH−(aq) produced (Table 1, eq 2). This titration is performed by adding a measured volume of KHP(aq) from a micropipet until the cathode compartment electrolyte is reverted back to the original colorless form (Figure S1C). Faraday’s laws of electrolysis state the amount of substances liberated at the electrodes of a cell is directly proportional to the quantity of electricity. The amount of different substances liberated by the same quantity of electricity depends upon the number of electrons involved in the respective electrode processes.3 The coefficients in the electrolytic and the chemical reactions (Table 1, eqs 1−6) represent the mole relation or the reaction stoichiometry. The electrolysis of water (eq 1) produces 2 mol of OH−(aq) in the cathode compartment. Passing the same charge in a concurrent anode compartment reaction (eq 3) produces 1 mol of I2(aq) as an electrolytic product of I−(aq). In the secondary chemical reactions (eqs 2 and 4), KHP(aq) and ascorbic acid react in one-to-one mole relation with the electrolytically produced OH−(aq) and I2(aq), respectively. Because the same current is flowing through the cathode and the anode, eqs 1−4 account for 1:1:0.5:0.5 mol relation between KHP(aq), OH−(aq), I2(aq), and ascorbic acid, respectively. The stoichiometry presented in these equations facilitates the direct determination of ascorbic acid from the amount of KHP(aq) added to the cathode compartment, independent of the magnitude of current flowing through the cell.

stoichiometry are presented in Table 1. An electrolysis cell (Figure 1), consisting of platinum electrodes and a salt bridge, is assembled. In the first experiment, a known mass of vitamin C powdered tablet sample, KI(aq), a buffer, and starch indicator are placed in the inner anode compartment. Electrolytic oxidation of KI(aq) produces I2(aq) with 100% current efficiency (i.e., without generating any side reaction products not effective in the chemical reaction of interest).9 As a direct current from a 9

Figure 1. Schematic diagram of the electrolysis cell: (a) 10 mL of 0.2 M KNO3(aq) and 1 drop of 0.2% phenolphthalein indicator, (b) fritted glass, (c) platinum wire cathode, (d) a switch, (e) Teflon holder for anode, (f) glass test tube, (g) sample, electrolyte, and indicator (vitamin C analysis, powdered sample, 100 mg of KI, 3 mL of 0.2 M acetate/acetic acid buffer, 2 drops of 2% starch; iron tablet analysis, powdered sample, 5 mL of 0.15 M Ce2(SO4)3(aq) in 1 M sulfuric acid, 1 drop of ferroin indicator), (h) platinum wire anode, (i) layer of agarKNO3 jelly, (j) stir bars (in both compartments). 899

dx.doi.org/10.1021/ed400630u | J. Chem. Educ. 2014, 91, 898−901

Journal of Chemical Education

Laboratory Experiment

Table 2. Summary of the Experimental Results for Vitamin C and Iron Supplement Tablets Analysis Vitamin C Supplement Tablets Iron Supplement Tablets a

Proposed Electrolytic Titration (average of 6 trials)

Proposed Electrolytic Titration from Slopes of Plots (Figures 2 and S2)

Volumetric Titration (average of six trials)

Manufacturer’s Label

0.224 g/tableta

0.227 g/tablet

0.249 g/tabletb

0.250 g/tablet

0.064 g/tabletc

0.063 g/tablet

0.0641 g/tabletd

0.065 g/tablet

Standard deviation ±4.5 × 10−3. bStandard deviation ±1.1 × 10−3. cStandard deviation ±1.5 × 10−3. dStandard deviation ±7.9 × 10−4.

determined from these trials was 0.224 g (standard deviation within 6 trials of titration ±4.5 × 10−3) (Table 2). The volume of KHP(aq) required to reach the end point in the cathode compartment linearly changed with the mass of the sample in the anode compartment (Figures 2 and S2). This

In the second experiment, the cathode compartment reactions (Table 1, eqs 1 and 2) are coupled with the electrolytic production of Ce4+(aq) and the chemical oxidation of Fe2+(aq) in the anode compartment (Table 1, eqs 5 and 6). A known mass of powdered iron supplement tablet, an excess amount of Ce3+(aq) in the acidic medium, and ferroin (used as an oxidation−reduction indicator) are placed in the inner anode compartment (Figure S1D). Electrolytic oxidation of Ce3+(aq) produces Ce4+(aq).10 Electrolytically produced Ce4+(aq) chemically oxidizes Fe2+(aq) to Fe3+(aq) and the end point of this titration is determined by a visual color change of the ferroin indicator from peach/red to pale blue (Figure S1E). The electrolysis is promptly stopped at the end point. The electrolytically produced OH−(aq) at the cathode and the presence of phenolphthalein indicator turn the electrolyte pink (Figure S1E). As stated earlier, the electrolytically produced OH−(aq) is titrated against KHP(aq) to determine the amount of OH−(aq) produced (Figure S1F). Equations, 2, 5, and 6 account for 1:1:1:1 mol relation between KHP(aq), OH−(aq), Ce4+(aq), and Fe2+(aq), respectively. The stoichiometry presented in these equations facilitates the direct determination of Fe2+(aq) from the amount of KHP(aq) added.

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EXPERIMENTAL DETAILS Two experiments presented in this paper are suitable for the undergraduate quantitative analysis laboratory. Students work in groups of three, and each experiment can be completed in 3 h. Each group performs 6 trials of the titration. Students obtain data so that they can analyze the mass of active ingredient in the dietary tablet using one of the analysis methods as indicated in Table 2, that is, either (1) direct calculations from electrolytic titration using similar masses of the powdered dietary tablet or (2) slope developed from electrolytic titration using different masses of the powdered dietary tablet.

Figure 2. A plot of the volume of 0.100 M KHP(aq) added to the cathode compartment versus the mass of vitamin C sample placed in the anode compartment. Coiled platinum wires served as electrodes and a layer of agar−KNO3 mixture and a fritted glass served as a salt bridge. The electrolyte temperature was 23 °C.

linear dependence confirmed the stoichiometry presented in the electrolytic and chemical reactions (eqs 1−6). Blank trials for each sample were conducted to confirm the negligible interference from electrolyte. The slope of the plot of the volume of KHP(aq) versus the mass of powdered vitamin C tablet was 90.4 (Figure 2). This slope indicated 90.4 μL of 0.100 M KHP(aq) neutralized in the cathode compartment was equivalent to 1 mg powdered vitamin C tablet in the anode compartment. The volume of KHP(aq) quantified from the slope corresponded to 9.04 μmol of KHP(aq) or 4.52 μmol concentration of ascorbic acid present in 1 mg powdered supplement tablet. Because the average mass of the tablet was 0.2846 g, one whole tablet contains 1.29 mmol or 0.227 g of ascorbic acid (Table 2). Both student authors (C.S. and K.G.) obtained the data independently. Similarly, the average mass of iron was determined to be 0.064 g (standard deviation within 6 trials of titration ±1.5 × 10−3 (Table 2). The slope of the plot of the volume of KHP(aq) versus the mass of powdered iron tablet was 29.4 (Figure S2). From this slope and the stoichiometry presented in eqs 1, 2, 5, and 6, 1.13 mmol Fe or 0.0631 g was present in a tablet.

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HAZARDS Approved safety goggles must be used while performing the experiment. Laboratory chemicals including potassium nitrate, potassium iodide, potassium hydrogen phthalate, potassium acetate, iron(II) sulfate heptahydrate, and cerium(III) sulfate can cause eye and skin irritation. Sulfuric acid can cause severe burns. Reagents involving acids must be prepared in a fume hood. H2(g) is produced at the cathode at an estimated rate of 0.1−0.2 mL/min (in a measured current range of 8−20 mA). Hydrogen gas is flammable, and a fume hood or a lecture room with a proper exhaust system must be used for a large size class. Instructors must examine and approve the electrode connections before students can begin the titrations. Appropriate waste containers must be used to collect the waste chemicals.

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RESULTS The electrolytic titrations were performed with approximately 0.0100 g of the sample. The average mass of ascorbic acid 900

dx.doi.org/10.1021/ed400630u | J. Chem. Educ. 2014, 91, 898−901

Journal of Chemical Education

Laboratory Experiment

Table 3. Summary of the Features of Proposed Electrolytic Products Titration, Coulometric Titrations, and Volumetric Titrations Proposed Electrolytic Products Titrations

Coulometric Titrations

•Reagent standardization is not required; less waste chemicals are produced •Powered by a 9 V battery (no electronic circuit)

•Reagent standardization is not required; less waste chemicals are produced •A constant current source or monitoring of charge is required •Covers the pedagogy of electrolysis and stoichiometry in redox reactions

•Covers the pedagogy of electrolysis and stoichiometry in redox reactions

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Both samples were independently analyzed by the conventional volumetric titration against standardized reagents.1 The results of the proposed method and volumetric titrations were in agreement with the manufacturer’s label. The experimental results are summarized in Table 2.

DISCUSSION The coulometric titrations3−7,11,12 and the electrolytic products’ titrations presented in this paper use costly precious metal electrodes. However, the proposed method offers added pedagogical value and other cost-effective features (Table 3). The electrolysis cell and titrations presented in this study are developed by an undergraduate student (C.S.) as a directed research project. Applications of the stoichiometric relations between the electrolytic products to the analysis of samples directly address the pedagogical objectives. The agreement between the results of the proposed electrolytic product titrations, volumetric titrations, and manufacturer’s label highlight the relevance of the experiments to the undergraduate quantitative analysis laboratory. Volumetric measurements tied to the stoichiometry of redox reactions and the electrolytically produced reagents indicate the significance of the experiments presented in this paper to the undergraduate quantitative analysis curriculum.13,14 ASSOCIATED CONTENT

S Supporting Information *

Directions to make the electrolysis cell, experimental procedure, and the pictures of the electrolysis cell. This material is available via the Internet at http://pubs.acs.org.

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•Covers the pedagogy of stoichiometry in redox reactions

REFERENCES

(1) Jeffery, G. H.; Bassett, J.; Mendham, J.; Denney, R. C. Vogel’s Textbook of Quantitative Chemical Analysis, 5th ed.; Longman: Essex, U.K., 1994; pp 381−393. (2) Day, R. A.; Underwood, A. L. Quantitative Analysis, 6th ed.; Prentice Hall: Englewood Cliffs, NJ, 1991; p 51. (3) Sawyer, D. T.; Sobkowiak, A.; Roberts, J. L. Electrochemistry for Chemists, 2nd ed.; Wiley: New York, 1995; pp 139−160. (4) Evans, D. H. Coulometric Titration of Cyclohexene with Bromine. J. Chem. Educ. 1968, 45 (2), 88−90. (5) Swim, J.; Earps, E.; Reed, L. M.; Paul, D. Constant-Current Coulometric Titration of Hydrochloric Acid. J. Chem. Educ. 1996, 73 (7), 679−683. (6) Lowinsohn, D.; Bertotti, M. Coulometric Titrations in Wine Samples: Studies on the Determination of S(IV) and the Formation of Adducts. J. Chem. Educ. 2002, 79, 103−105. (7) Dabke, R. B.; Gebeyehu, Z.; Petermann, M.; Johnson, N., Jr.; Patel, K. Analysis of Household Products: Coulometric Titration Experiment in the Undergraduate Laboratory. Chem. Educ. 2011, 16, 160−163. (8) Jeffery, G. H.; Bassett, J.; Mendham, J.; Denney, R. C. Vogel’s Textbook of Quantitative Chemical Analysis, 5th ed.; Longman: Essex, U.K., 1994; pp 261−262. (9) Lingane, J. J. Electroanalytical Chemistry, 2nd ed.; Interscience: New York, 1958; pp 551−552. (10) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001; p 432. (11) Williams, K. R.; Young, V. Y.; Killian, B. J. Coulometric Titration of Ethylenediaminetetraacetate (EDTA) with Spectrophotometric Endpoint Detection: An Experiment for the Instrumental Analysis Laboratory. J. Chem. Educ. 2011, 88, 315−316. (12) Dabke, R. B.; Gebeyehu, Z.; Thor, R. Coulometric Analysis Experiment for the Undergraduate Chemistry Laboratory. J. Chem. Educ. 2011, 88, 1707−1710. (13) Harvey, D. Modern Analytical Chemistry; McGraw Hill: Boston, MA, 2000; pp 273−367. (14) Skoog, D. A.; Holler, F. J.; Crouch, S. R. Principles of Instrumental Analysis, 6th ed.; Thomson Higher Education: Belmont, CA, 2007; pp 707−712.

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Volumetric Titrations •Reagent standardization is required; more waste chemicals are produced •No power source is required

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 (Columbus State University, Columbus, GA) and David L. Pringle (University of Northern Colorado, Greeley, CO) for their helpful comments and suggestions. We thank the reviewers of the manuscript for their valuable suggestions. Christopher Scanlon gave a poster presentation of this work at the 245th ACS National Meeting and Exposition, New Orleans, LA in the spring 2013. We gratefully acknowledge Tim Howard (Columbus State University, Columbus, GA) for supporting Kameron’s summer research under the National Science Foundation’s Robert Noyce Teacher Scholarship Program (Grant 1136356). 901

dx.doi.org/10.1021/ed400630u | J. Chem. Educ. 2014, 91, 898−901