Field and In-Lab Determination of Ca2+ in Seawater - Journal of

Laboratory Experiment pubs.acs.org/jchemeduc

Field and In-Lab Determination of Ca2+ in Seawater Robin Stoodley,* Jose R. Rodriguez Nuñez, and Tessa Bartz Department of Chemistry, University of British Columbia, Vancouver, British Columbia V6T 1Z1, Canada S Supporting Information *

ABSTRACT: Portions of classic undergraduate quantitative analysis experiments in complexiometric titration and potentiometry are combined with a field-sampling experience to create a two period (2 × 3 h) comparison-based experiment for second-year students. A multifunctional chemical analysis device is used with calcium ion-selective electrode for field measurement of calcium concentration in seawater. The results are compared with later simultaneous potentiometric and colorimetric titration with ethylenediaminetetraacetic acid (EDTA). Students compare precision, difficulty, and sources of error of the different methods and present a written argument supporting their choice of best method including advantages and limitations. The experiment provides an introduction to chemistry in-the-field, sample collection and handling, complex matrices, imperfect methods, fallibility of instrumentation, real-world performance, and decision-making. KEYWORDS: Second-Year Undergraduate, Analytical Chemistry, Laboratory Instruction, Hands-On Learning/Manipulatives, Problem Solving/Decision Making, Quantitative Analysis, Titration/Volumetric Analysis, Water/Water Chemistry theme among many students’ first quantitative analysis experiences is the determination of sea- or surface water by complexiometric titration.1−4 Sometimes presented as a measurement of water hardness, it is typical to measure Ca2+ concentration or the sum of Ca2+ and Mg2+ concentrations. This theme is common because it shows application relevance and offers overlap with typical lecture concepts: complex matrices, simultaneous equilibria, visual indicators, end point versus equivalence point, accuracy and precision, and so forth. Many universities offer laboratory experiments on potentiometry in the same course; determination of fluoride in toothpaste is classic,5 other examples include chloride anion in natural waters6 or in aquariums,7 indirect determination of reducing sugars8 via Cu2+ or of carbon dioxide.9 Experiments of this type support concepts of calibration curve and standard addition methods, the Nicolsky-Eisenman equation, ionic strength, matrix matching, and accuracy and precision, and so forth. At the University of British Columbia, the experiments of ethylenediaminetetraacetic acid (EDTA) titration of synthetic seawater and potentiometric determination of fluoride in toothpaste have been run for many years. As Mossman et al.4 said “there is a certain staleness to the established canon of quant experiments”; he remedied this by putting a twist on the classic EDTA-Ca titration by incorporating photometric end point detection. Offered here is a twist on two classic experiments by combining their learning objectives and adding an in-field component. The experiment covers most of the concepts above through comparison of in-field direct potentiometric measurement of Ca2+ concentration in seawater with potentiometric and colorimetric titration in-lab. A Vernier (Beaverton, OR) LabQuest multifunctional chemical analysis (MCA)10 device with a calcium ion-selective electrode is used

A

© XXXX American Chemical Society and Division of Chemical Education, Inc.

for potentiometry. Through the in-field part of the experiment, sampling theory and practice is introduced. Relatively few examples of experiments including in-field sampling exist in the literature,11−14 yet it is a topic of considerable importance. ACS guidelines15 suggest that the student experience avoid overemphasizing measurement and instead reflect the full analytical process including obtaining the sample. In revising our laboratory curriculum to include a comparison of methods, we sought to engage students at a higher cognitive level than in our classic experiments. Domin16 proposed that traditional lab activities are designed for loworder cognitive skills instead of high-order ones such as analysis, synthesis, and evaluation, as defined by Bloom’s taxonomy of educational objectives.17 In “cookbook” or expository18 experiments, students often do not need to consider the characteristics of the sample, as they have been prespecified. The experimental method has been predetermined, and its limitations are often not apparent because the sample has been selected to suit the instrument. While parts of this experiment are expository, the comparison of results from the three methods challenges the students to critically evaluate their data and decide which method is preferred. Literature suggests that comparison activities can lead to greater learning gains than traditional instruction.19

■

EXPERIMENTAL OVERVIEW Students work in pairs for the duration of the experiment. In the first lab period, students prepare their field trip kits including the LabQuest device, electrode, low and high

A

dx.doi.org/10.1021/ed4005722 | J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education

Laboratory Experiment

the course. The experiment is completed over two, 3-h laboratory periods spaced 1 week apart.

concentration calibration solutions, and sample collection bottle. Accompanied by a teaching assistant, students walk to nearby Wreck Beach, located on the delta of the Fraser River where it meets the Pacific Ocean. The beach seawater is naturally diluted by the incoming riverwater. The extent of dilution varies according to river flow rate, weather, and tides; 40% variation in Ca2+ concentration has been observed within one tidal cycle. Student determined calcium concentration cannot be compared to any fixed value and are expected to be lower than that of standard composition seawater. At the beach, the teaching assistant describes best practices in sample collection. For safety, students do not enter the water but instead take grab samples from a rock groin. Students filter the seawater into their sample bottle and transfer a second portion to a plastic beaker for immediate measurement. Students calibrate the instrument with low (10 mg/L Ca2+) and high (1000 mg/L Ca2+) standard solutions before measuring the seawater. Temperature of each solution is recorded via a Vernier thermocouple probe. Ca2+ concentration versus time is measured for 180 s with six replicate measurements. Students return to the laboratory to analyze their data and extract six values for Ca2+ concentration from this “direct potentiometric” method. The collected seawater samples are stored in the refrigerator until the following lab period 1 week later. In the second period, students conduct simultaneous colorimetric and potentiometric titration of their seawater, giving students experience in both visual and quantitative end point detection. Colorimetric titration of metal ions is well covered in textbooks and literature.20−23 EDTA is used as a complexing agent, Calcon (1-(2-hydroxy-1-naphthylazo)-2naphthol-4-sulfonic acid sodium salt) is used as a colorimetric indicator, and diethylamine is used to raise pH to 12.5 to precipitate Mg as Mg(OH)2. Left in solution, the Mg would otherwise also bind EDTA. Potentiometric titration is carried out with the LabQuest and calcium electrode. Volume of EDTA added and potentiometric Ca2+ concentration reading are recorded into the lab notebook. Students choose how to determine the end point from their data; first and second derivative methods are common choices. Students repeat the titration in triplicate. During the second lab period, students are given an oral quiz, which asks the open-ended question: identify and rationalize which method is best and which worst. Most students select the potentiometric titration as best and mention the poor accuracy and precision of the direct potentiometry or the precision of the colorimetric titration. In their lab reports, students compare their determined Ca2+ concentration by the three different methods, including use of statistical comparison via Student’s t test. For most student measurements, the direct potentiometric result differs from both titration results, although the titration results are normally in statistical agreement. After completing the calculations and having had modest prompting about the potential for bias in each method, students typically identify in their reports the colorimetric titration as most accurate and usually most precise. Additional experimental details including specifics of discussion questions about method comparison and end point detection are included as Supporting Information. This experiment has been successfully run for three terms of our second-year quantitative analytical chemistry course in the 2012 and 2013 academic years. A total of 344 students were enrolled. Majors in chemistry, environmental science, medical lab science, pharmacology and physiology are required to take

■

HAZARDS Diethylamine is highly corrosive, flammable, and has an unpleasant odor. Skin contact may result in chemical burns. Eye protection, lab coat, and gloves must be worn. A bottle-top dispenser in the fumehood is used to minimize potential contact. When collecting the seawater samples and conducting potentiometry at the public beach, we have opted for students not to wear labcoats and eye protection to avoid alarming members of the public. To avoid breakage, no glassware is used at the beach. Students are accompanied at the beach by a teaching assistant who carries a first aid kit.

■

RESULTS AND DISCUSSION Representative data is drawn from all laboratory sections offered in spring 2013. Students worked in pairs for the duration of the experiment, giving 54 independent analyses. Each pair of students calibrated their own calcium ion-selective electrode, collected their own sample, operated the LabQuest with ion-selective electrode, and titrated their sample. Figure 1 shows representative student titration data for a sample of seawater. Readout of the LabQuest as the titration proceeds is given in panel A. Calcon addition lowers the Ca 2+

Figure 1. Representative student titration data. (A) LabQuest readout of calcium ion-selective electrode. Electrode was calibrated immediately before measurement. The effect of Calcon and diethylamine addition to the diluted seawater (SW) are shown. (B) Natural logarithm of LabQuest Ca readout, which returns the data to a form equivalent to electrode voltage. Colorimetric end point was found when 10.45 mL of EDTA was added. (C) Negative of the first derivative of the plot in (B). The negative is taken so the y values of data are positive. End point was found by first derivative method to be when 10.14 mL of EDTA was added. B




Table 1. Summary of 2013 Spring Semester Student Resultsa Method

Averageb [Ca2+]/ mg/L

Average SD/ mg/L

% Average RSD

Maximum determined [Ca2+] / mg/L

Minimum determined [Ca2+]/ mg/L

Direct potentiometry Potentiometric titration Colorimetric titration

153 261 280

14.8 5.56 5.30

9.72 2.16 1.85

897 322 410

4.68 162 175

a Average calcium concentration found by different methods, the average over each student’s standard deviation (SD), the average relative standard deviation (RSD) and range of results (maximum and minimum) are given. bMeasurements were not conducted on replicate samples; field measurements were made 12 times over a 4 week period. Average is given here to allow intermethod comparison of result and relative precision.

disadvantages of the different analysis methods. A student comment demonstrates the value: “The lab challenged my instinctive assumption that quantitative measurements (e.g., with the electrode) are always better than qualitative (e.g., the color change). It prompted me to think twice about what was really going-on chemically.”

concentration reading, although subsequent addition of diethylamine partially returns the value. The natural logarithm of electrode readings are shown in panel B, and to demonstrate the most common method students use to find the end point, the negative first derivative of the panel B data is shown in panel C. The potentiometric end point corresponds to 203 mg Ca2+/L and the colorimetric end point corresponds to 209 mg Ca2+/L. Instructor analysis of this sample by atomic absorption spectrophotometry using a standard addition method yielded Ca2+ concentration of 211 ± 3 mg/L (95% confidence level). Results from direct potentiometry (Table 1) show poor accuracy and precision (high relative standard deviation). Even under ideal conditions, the Vernier electrodes give poor accuracy and reproducibility compared to Thermo Scientific Orion calcium electrodes (part# 9320BN), but there are other potential causes for the poor results. The voltage output from an electrode selective for an ion i, in the absence of interference, varies with the activity (not the concentration) of the ion: RT E = constant + ln ai ziF

■

ASSOCIATED CONTENT

* Supporting Information S

A lab procedure for students and a series of notes for the instructor are available. This material is available via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS R.S. thanks our department’s ChIRP Teaching and Learning Enhancement Fund for a grant permitting purchase of the LabQuest devices and ion-selective electrodes.

(1)

where E is the measured voltage, R is the universal gas constant, T is the sample temperature in Kelvin, zi is the charge on the ion of interest, F is Faraday’s constant, and ai is the activity of i. Activity varies with ionic strength, complicating the direct potentiometry: the calibration standards should have equal ionic strength as the sample. Variation in riverwater dilution of the seawater means the correct ionic strength for our calibration standards can only be approximated. NaCl is added to reach a total ionic strength of 0.45 M. Similarly, the temperature of the standard solutions and sample should also be matched to avoid error. The variation between maximum and minimum student titration results is indicative of the natural variation of Ca2+ concentration in the seawater during the term. In comparing results of the colorimetric and potentiometric titrations a systematic bias toward lower concentration for the latter was observed. Added Calcon indicator binds a portion of the total calcium and masks it from potentiometric detection. Near the end point, Calcon is displaced from calcium by EDTA, but the calcium remains masked from the electrode. The added Calcon (∼4 × 10−6 mol) is about 3−5% of the typical amount of calcium, in agreement with the magnitude of the bias.

■

REFERENCES

(1) Yang, S.; Li, C. Using Student-Developed, Inquiry-Based Experiments To Investigate the Contributions of Ca and Mg to Water Hardness. J. Chem. Educ. 2009, 86, 506−513. (2) Selco, J.; Roberts, J.; Wacks, D. The Analysis of Seawater: A Laboratory-Centered Learning Project in General Chemistry. J. Chem. Educ. 2003, 80, 54−57. (3) Belle-Oudry, D. Quantitative Analysis of Sulfate in Water by Indirect EDTA Titration. J. Chem. Educ. 2008, 85, 1269−1270. (4) Mossman, D.; Kooser, R.; Welch, L. The Complexometric Determination of Calcium and Magnesium in Limestone Using a Laser Photometer for Endpoint Identification. J. Chem. Educ. 1996, 73, 82− 85. (5) Light, T.; Cappuccino, C. Determination of Fluoride in Toothpaste Using an Ion-Selective Electrode. J. Chem. Educ. 1975, 52, 247−250. (6) Berger, M. Potentiometric Determination of Chloride in Natural Waters: An Extended Analysis. J. Chem. Educ. 2012, 89, 812−813. (7) Harris, T. Potentiometric Measurements in a Fresh-Water Aquarium. J. Chem. Educ. 1993, 70, 340−341. (8) Moresco, H.; Sanson, P.; Seoane, G. Simple Potentiometric Determination of Reducing Sugars. J. Chem. Educ. 2008, 85, 1091− 1093. (9) Kocmur, S.; Corton, E.; Haim, L.; Locascio, G.; Galagosky, L. CO2-Potentiometric Determination and Electrode Construction, a Hands-on Approach. J. Chem. Educ. 1999, 76, 1253−1255. (10) Vannatta, M. W.; Richards-Babb, M.; Solomon, S. D. Personal Multifunctional Chemical Analysis Systems for Undergraduate Chemistry Laboratory Curricula. J. Chem. Educ. 2010, 87, 770−772.

■

CONCLUSION Building from two classic “quant” experiments, this experiment develops a new learning experience that exposes students to field chemistry. Students engaged more fully in this experiment compared to our traditional experiments. Students worked with complex, real-world data and demonstrated high-order cognitive skills while critically considering the advantages and C




(11) Jeannot, M. Analysis of iron in lawn fertilizer: A sampling study. J. Chem. Educ. 2006, 83, 243−244. (12) Luck, L. A.; Blondo, R. M. The Grapes of Class: Teaching Chemistry Concepts at a Winery. J. Chem. Educ. 2012, 89, 1264−1266. (13) Sinniah, K.; Piers, K. Ion chromatography: Analysis of Ions in Pond Waters. J. Chem. Educ. 2001, 78, 358−362. (14) Bachofer, S. J. Sampling the Soils around a Residence Containing Lead-Based Paints: An X-ray Fluorescence Experiment. J. Chem. Educ. 2008, 85, 980−982. (15) American Chemical Society Committee on Professional Training, Analytical Chemistry Supplement 2008, CTP005610. (16) Domin, D. A Content Analysis of General Chemistry Laboratory Manuals for Evidence of Higher-Order Cognitive tasks. J. Chem. Educ. 1999, 76, 109−112. (17) Bloom, B. S., Ed. In Taxonomy of Educational Objectives, the Classification of Educational GoalsHandbook I: Cognitive Domains; David McKay Co. Inc: New York, NY, 1956. (18) Domin, D. A Review of Laboratory Instruction Styles. J. Chem. Educ. 1999, 76, 543−547. (19) Alfieri, L.; Nokes-Malach, T. J.; Schunn, C. D. Learning through Case Comparisons: A Meta-Analytic Review. Educ. Psychol. 2013, 48 (2), 87−113. (20) Day, R. A.; Underwood, A. L. Quantitative Analysis, 6th ed.; Prentice-Hall: Upper Saddle River, NJ, 1991; p 622. (21) Hage, D. S.; Carr, J. D. Analytical Chemistry and Quantitative Analysis; Pearson Prentice Hall: Upper Saddle River, NJ, 2011; pp 319−341. (22) Olsen, K.; Ulicny, L. Reduction of Calcium Concentrations by the Brita (R) Water Filtration System: A Practical Experiment in Titrimetry and Atomic Absorption Spectroscopy. J. Chem. Educ. 2001, 78, 941−941. (23) Yappert, M.; DuPre, D. Complexometric Titrations: Competition of Complexing Agents in the Determination of Water Hardness with EDTA. J. Chem. Educ. 1997, 74, 1422−1423.

D


Field and In-Lab Determination of Ca2+ in Seawater - Journal of

Recommend Documents