Communication Cite This: J. Chem. Educ. 2019, 96, 1283−1286
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BYOL: Bring Your Own Lime Hands-On Laboratory Experience Mikhail Kurushkin,* Chantal Tracey, and Maria Mikhaylenko Chemistry Education Research and Practice Laboratory, SCAMT Institute, ITMO University, 9 Lomonosova Street, Saint Petersburg 191002, Russian Federation
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S Supporting Information *
ABSTRACT: BYOL (Bring Your Own Lime), a novel laboratory experiment, is introduced in this paper. Students are encouraged to bring affordable household and grocery store chemicals into the classroom. The Bring Your Own Chemical (BYOC) concept demonstrates to students that food and household products are chemicals too. Citruses are often used to highlight the link between everyday life and chemistry. Due to the large amount of citric acid that limes contain, they can be squeezed and their juice used for experimentation. The aim of the experiment is to determine the first stepwise dissociation constant of citric acid in grocery store limes. The BYOL: Bring Your Own Lime laboratory experiment achieves several goals: it makes chemistry relatable, boosts both content knowledge (preparative chemistry, spectrophotometry and conductivity measurements) and soft skills, is not time-consuming (requires 75 min if the provided calibration curve is used), and is an enjoyable experience. KEYWORDS: First-Year Undergraduate/General, Collaborative/Cooperative Learning, Hands-On Learning/Manipulatives, Misconceptions/Discrepant Events, Acids/Bases, Aqueous Solution Chemistry, Conductivity, Laboratory Equipment/Apparatus, UV−Vis Spectroscopy
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ecently, BYOD (Bring Your Own Device) has been becoming a global trend in science education. This is unsurprising as it has the advantage of making science relatable to the real world. It ignites interest in students and promotes teamwork during lessons.1−3 Despite these obvious advantages, only one article is dedicated to BYOD in chemical education: It describes using a digital notebook in a classroom setting.4 Another promising concept that has been initiated is the “build your own device” concept, introduced in a BYOP (Build Your Own Photometer) article.5 Overall, it is safe to say that this technique has been overlooked and rather underutilized, regardless of the undoubtable educational benefits associated with it. BYOL (Bring Your Own Lime), a novel laboratory experiment, is introduced in this paper. As its name implies, students are encouraged to bring their own chemicals to class. These chemicals take the form of affordable household and/or grocery store items. Advantages of Bring Your Own Chemical (BYOC) approaches include easier and less time-consuming laboratory preparations, if any, resulting in higher overall efficiency of the experiment itself. Moreover, the BYOC idea demonstrates to students that food and household products are chemicals too, dispelling the common misconception that chemistry is abstract and distinct from everyday life.6,7 BYOC also helps students engage with chemical concepts through collaborative and hands-on learning. BYOC therefore makes chemistry relatable and relevant to students, piquing their interest in the science. © 2019 American Chemical Society and Division of Chemical Education, Inc.
PREVIOUS EFFORTS
Two laboratory experiments involving household vinegar have been published.8,9 The first reported experiment is dedicated to determining the percentage of acetic acid in household vinegar, while the second affords students the opportunity to study the reaction between acetic acid and sodium bicarbonate, simple baking soda. Citruses have also been used in laboratory exercises.10,11 One experiment is aimed at the production of liqueur from lemons, while the other uses citrus peels in the reduction of Cu(II). Vinegar and citruses seem to be rather common and laboratory-friendly household chemicals. Vinegar, however, has a number of safety issues associated with its use, being hazardous in the case of skin or eye contact, inhalation, and ingestion, and it may lead to tissue damage. Furthermore, vinegar’s scent is undeniably unpleasant and can cause headaches and create a negative atmosphere in the classroom. Therefore, using citruses for experimentation is preferable. Due to the large amount of citric acid that limes and lemons contain, they can be squeezed and their juices collected. The first stepwise dissociation constant of citric acid contained within these fruits is ultimately determined, which is the aim of the experiment. pH and conductivity measurements12 can be Received: November 22, 2018 Revised: April 10, 2019 Published: April 30, 2019 1283
DOI: 10.1021/acs.jchemed.8b00966 J. Chem. Educ. 2019, 96, 1283−1286
Journal of Chemical Education
Communication
Λc 1 1 = Λ Λ° K a(Λ°)2
used to determine the dissociation constant. As the latter is more credible,12 that is the method that will be employed. The use of household and grocery store chemicals, prepared by the teacher beforehand, has been previously introduced into laboratory practice. In this experiment, however, the students carried out the method themselves, in pairs, from start to finish, promoting the discovery and learning aspects of the activity. It is of equal importance that the students employed a combination of techniques for this new experiment: liquid aliquot method, UV spectrophotometry, and conductivity measurements. All of these are indispensable for first-year undergraduates.
Here, Λ° is the molar conductivity at infinite dilution (conductivity at cΛ = 0) in S m2 mol−1. The molar conductivity at infinite dilution for the temperature range from 290 to 300 K is given in Table 1.19 Table 1. Molar Conductivity at Infinite Dilutiona T/°C
T/K
Λ°/S m2 mol−1
17 18 19 20 21 22 23 24 25 26 27
290 291 292 293 294 295 296 297 298 299 300
0.0370 0.0377 0.0383 0.0390 0.0396 0.0402 0.0409 0.0415 0.0422 0.0428 0.0435
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THEORETICAL BACKGROUND In an aqueous solution of citric acid homogeneous equilibrium establishes H3Cit(aq) ↔ H(aq)+ + H 2Cit(aq)−
Polyprotic acids (e.g., citric acid which has triprotic nature) dissociate stepwise. The above-mentioned reaction describes the first step of dissociation of citric acid. The quantitative measure of homogeneous equilibrium is the dissociation constant (Ka): K a1 =
a
[H(aq)+ ][H 2Cit(aq)− ] [H3Cit(aq)]
Inferred from ref 19.
The degree of dissociation (α) given by Λ Λ° and combined with the Ostwald dilution law gives the equation used to calculate the first stepwise dissociation constant: α=
Thus, the above-mentioned equation describes the first stepwise dissociation constant of citric acid. For citric acid, Ka1 = 7.4 × 10−4, Ka2 = 1.7 × 10−5, and Ka3 = 4.0 × 10−7.13 Due to the second and third steps being less considerable, it is acceptable to neglect them. While Ka is a quantitative measure, the classification into strong and weak electrolytes is conventional. The strength of a potential electrolyte is characterized by its degree of dissociation in a solution. However, degree of dissociation (the fraction of molecules that underwent dissociation) depends on molar concentration. Hence, for aqueous solutions of acids, hydronium ion H3O+ (pKa = −1.74) is well-used as an approximation of the state of protons in water.14 Those acids with pKa > −1.74 (Ka < 101.74) are considered weak. A few methods15−17 exist in which the concentration of an aqueous acid can be determined, but as accuracy is key, spectrophotometry gives the most accurate results. With the preparation of a series of citric acid solutions of known concentrations and measurements of each absorbance, a calibration curve is obtained. The optimal wavelength for measuring absorbance, λmax, for citric acid is 209 nm.18 This falls within the ultraviolet (UV) region of the spectrum. In UV spectrophotometry, really dilute solutions, on the order of 10−4−10−6, are required. The electrical conductivity (σ) (often called specific conductance in older literature) is the reciprocal of the resistivity and is usually given in mS cm−1 or μS cm−1. The molar conductivity (Λ) of an electrolyte solution is defined as the conductivity divided by amount-of-substance concentration and is given in S m2 mol−1: σ Λ= c
α 2c 1−α Traditionally, for weak electrolytes (α ≪ 1) the degree of dissociation is omitted in the denominator making it equal to 1. For more accurate results, it is not advised. Ka =
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EXPERIMENTAL OVERVIEW In total, 87 students participated in the BYOL activity. Six groups of approximately 14 students each were asked to take their own limes to the laboratory and given a copy of the laboratory guide (see Supporting Information) beforehand. On the day of the experiment, the students’ excitement was palpable. The students worked in pairs. Each pair was first required to determine the citric acid concentration of their lime juice by using UV spectrophotometry. They interpolated their values from the calibration curve provided in Figure 1 (see Supporting Information for absorbances). Each pair of students then made a solution series of lime juice with varying concentrations by performing simple dilutions. The conductivity of each solution was measured using a conductivity meter and each temperature noted. The students were then given a post-test (see Supporting Information) immediately after the experiment. The activity and the post-test took each group on average 75 min to complete. Students were required to hand in a completed laboratory report within 2 weeks.
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HAZARDS It is necessary to wear safety glasses during the experiment due to possible eye irritation in case of contact with citric acid or
The Ostwald dilution law is used to express the concentration dependence of the molar conductivity of an electrolyte solution: 1284
DOI: 10.1021/acs.jchemed.8b00966 J. Chem. Educ. 2019, 96, 1283−1286
Journal of Chemical Education
Communication
Figure 2. Degree of dissociation as a function of concentration. Figure 1. Calibration curve. Green dashed line represents limes.
completed. A score of 50% was taken as the pass mark, and 56 of the 87 students (64%) were able to achieve this score. There were 33 students (38%) who were able to define the first stepwise dissociation constant of citric acid and were able to identify the relationship between the concentration and degree of dissociation. There were 59 (68%) students who understood why it was necessary to determine the citric acid concentration of their lime juice while 57 (65%) recognized that the ambient temperature influenced the measured conductivity. There were 67 students (77%) who were able to state at least one factor that could have contributed to the difference between their calculated first stepwise dissociation constant and the literature value. These results suggest that the students did self-learning before performing the experiment, indicating that asking them to bring their own limes and providing them with the method prior to the activity piqued their interest.
lime juice. The laboratory instructor or the teacher should halve the limes themselves so that students do not use or bring any sharp objects to the lab.
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RESULTS AND DISCUSSION Table 2 shows an example of the experimental results obtained. A discrepancy is observed between the experimentally obtained first stepwise dissociation constant of citric acid and the accepted reference value (7.4 × 10−4). This may be attributed to any instrumental errors incurred during the preparation of the solutions for the calibration curve and for the conductivity measurements. It should also be kept in mind that molar conductivity at infinite dilution has significant dependence on the ambient temperature. In order to validate the results, the degree of dissociation was plotted against concentration for the lime juice (Figure 2). The smooth curve suggests that the experiment provides accurate results and was conducted with precision. A pair of students was considered to have performed satisfactorily if their calculated constant lay between 6.4 × 10−4 and 8.4 × 10−4. According to the 42 laboratory reports provided by the students, 28 (67%) exhibited acceptable values of the first stepwise dissociation constant which signifies the fact that a majority of students were able to advance their laboratory skills working with the equipment and methods they had been unfamiliar with prior to the experience. The students were given a post-test. The aim of the test was to measure their understanding of the activity they had just
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CONCLUSION Measuring the conductivity has proven itself to be a valid and accurate way of determining the first stepwise dissociation constant of citric acid in lime juice, as evidenced by the closeness of 67% of the experimental values to the true value. The BYOC concept and the pleasant scent of the freshly squeezed citruses evoked anticipation for the day of experiment and resulted in a relaxed and fun atmosphere in the laboratory, which is a favorable prerequisite for successful learning. A combination of collaborative and hands-on learning techniques, in turn, provided students with an opportunity to advance the following skill set: preparative chemistry,
Table 2. Experimental and Calculated Results from Lime Juice with T = 294 K c/mol L−1
σ/μS cm−1
σ/S m−1
Λ/S m2 mol−1
Λ°/S m2 mol−1
α
0.1721 0.1583 0.1239 0.0895 0.0688 0.0413
4487 4285 3820 3251 2825 2170
0.4487 0.4285 0.3820 0.3251 0.2825 0.2170
0.0026 0.0027 0.0031 0.0036 0.0041 0.0053
0.0396 0.0396 0.0396 0.0396 0.0396 0.0396
0.0659 0.0683 0.0779 0.0917 0.1037 0.1327
Ka 8.0 7.9 8.1 8.3 8.3 8.4 8.2
× × × × × × ×
10−4 10−4 10−4 10−4 10−4 10−4 10−4 ⟨Ka⟩a
a
Arithmetic mean. 1285
DOI: 10.1021/acs.jchemed.8b00966 J. Chem. Educ. 2019, 96, 1283−1286
Journal of Chemical Education
Communication
(9) Lanni, L. M. Filling a Plastic Bag with Carbon Dioxide: A Student-Designed Guided-Inquiry Lab for Advanced Placement and College Chemistry Courses. J. Chem. Educ. 2014, 91 (9), 1390−1392. (10) Mendes, D. E.; Schoffstall, A. M. Citrus Peel Additives for OnePot Triazole Formation by Decarboxylation, Nucleophilic Substitution, and Azide-Alkyne Cycloaddition Reactions. J. Chem. Educ. 2011, 88 (11), 1582−1585. (11) Naviglio, D.; Montesano, D.; Gallo, M. Laboratory Production of Lemon Liqueur (Limoncello) by Conventional Maceration and a Two-Syringe System to Illustrate Rapid Solid-Liquid Dynamic Extraction. J. Chem. Educ. 2015, 92 (5), 911−915. (12) Nyasulu, F.; Moehring, M.; Arthasery, P.; Barlag, R. Ka and Kb from PH and Conductivity Measurements: A General Chemistry Laboratory Exercise. J. Chem. Educ. 2011, 88 (5), 640−642. (13) Lide, D. R. Handbook of Chemistry and Physics, CD-ROM Version 2010, 90th ed.; CRC Press/Taylor and Francis: Boca Raton, FL, 2010. (14) Jolly, W. L. Modern Inorganic Chemistry; McGraw-Hill, Inc.: New York, 1984. (15) Farajzadeh, M. A.; Nagizadeh, S. Citric Acid Determination by Dual Wavelength Spectrophotometry. J. Chin. Chem. Soc. 2002, 49 (4), 619−624. (16) Penniston, K. L.; Nakada, S. Y.; Holmes, R. P.; Assimos, D. G. Quantitative Assessment of Citric Acid in Lemon Juice, Lime Juice, and Commercially-Available Fruit Juice Products. J. Endourol. 2008, 22 (3), 567−570. (17) Umali, A. P.; Anslyn, E. V.; Wright, A. T.; Blieden, C. R.; Smith, C. K.; Tian, T.; Truong, J. A.; Crumm, C. E.; Garcia, J. E.; Lee, S.; et al. Analysis of Citric Acid in Beverages: Use of an Indicator Displacement Assay. J. Chem. Educ. 2010, 87 (8), 832−835. (18) Krukowski, S.; Karasiewicz, M.; Kolodziejski, W. Convenient UV-Spectrophotometric Determination of Citrates in Aqueous Solutions with Applications in the Pharmaceutical Analysis of Oral Electrolyte Formulations. J. Food Drug Anal. 2017, 25 (3), 717−722. (19) Apelblat, A.; Barthel, J. Conductance Studies on Aqueous Citric Acid. Zeitschrift fur Naturforsch. - Sect. A J. Phys. Sci. 1991, 46 (1−2), 131−140.
spectrophotometry, and conductivity measurements. When asked to comment on the activity, the students said it was satisfying to achieve a real result and that it was interesting to look at household objects in a chemical way. In conclusion, the Bring Your Own Lime laboratory experiment achieves several goals: makes chemistry relatable, boosts both content knowledge and soft skills, is not timeconsuming (requires 75 min if the provided calibration curve is used), and is an enjoyable experience.
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.8b00966.
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Instructor notes (PDF, DOCX) Instructor worksheet and calibration curve data (XLSX) Numberical problems (PDF, DOCX) Post-test answer key (PDF, DOCX) Post-test (PDF, DOCX) Student handout (PDF, DOCX)
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Mikhail Kurushkin: 0000-0001-9031-8247 Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS We greatly acknowledge all the students that participated in BYOL and gave feedback, making this work possible. REFERENCES
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DOI: 10.1021/acs.jchemed.8b00966 J. Chem. Educ. 2019, 96, 1283−1286
