Laboratory Experiment Cite This: J. Chem. Educ. XXXX, XXX, XXX−XXX
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Physicians as the First Analytical Chemists: Gall Nut Extract Determination of Iron Ion (Fe2+) Concentration Mary T. van Opstal,*,‡ Philip Nahlik,† Patrick L. Daubenmire,† and Alanah Fitch† †
Department of Chemistry and Biochemistry, Loyola University Chicago, 1032 West Sheridan Road, Chicago, Illinois 60660, United States ‡ Department of Chemistry, Harper College, 1200 West Algonquin Road, Palatine, Illinois 60067, United States S Supporting Information *
ABSTRACT: Many students enroll in college general chemistry with an interest in a medical career. In those (and alternative) careers, they will need to make critical decisions about data and how that data are acquired. A significant portion of introductory lab experiments are, in principle, but not necessarily in practice, devoted to understanding how chemical information is gained and how it is deemed reliable. Here we report on a laboratory experiment that is presented as a guided inquiry investigation grounded in the history of science and the link between early medicine and quantitative data. Oak gall extract is used to quantitatively determine the amount of iron present in a water sample using UV−vis absorbance. In this experiment, students are introduced to the Beer−Lambert law, calibration curves, and the reading of UV−vis spectra KEYWORDS: First-Year Undergraduate/General, History/Philosophy, Laboratory Instruction, Collaborative/Cooperative Learning, Inquiry-Based/Discovery Learning, UV−Vis Spectroscopy, Quantitative Analysis
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chemistry classroom activities.8−11 The lab experience presented here brings sensory and intuitive observation (color perception, derivation of chemical materials directly from nature) and quantitative observation (calibration curve with instrumentation) into one unified historically contextualized laboratory. There is overlap between the mathematical models of science and the sensory intuitive driven12−14 exploration of the NOS dominating scientific thought up to the advent of instrumentation. As cogently observed by Mamlok-Naaman et al.:15 If indeed, one can claim that children’s understanding of scientific concepts develops in a way analogous to that of the knowledge of scientists during the known history [intuitive and sensory], then presenting students with the steps of this development may help them “grow up” with science. This historical lab where students use the “First Chemical Reagent”16 to determine iron in water samples may help
INTRODUCTION The laboratory is a place for students to investigate, learn the processes of observation and science, make claims based on evidence, evaluate, and develop explanations and solutions to scientific problems.1 Students need to build these practices to gain skills for their future careers in science and engineering. Learning laboratories that move away from procedural “cookbook” experiments to guided inquiry and often openended experiments can better allow students to gain these skills. Guided inquiry laboratory instruction affords students the opportunities to use critical thinking skills,2 problem solving skills,3 metacognitive strategies,4 and writing skills.5 In addition to these skills, we hope to engage students with a better understanding of how science works, the nature of science (NOS).6 In addition to inquiry-based lab instruction, a historical perspective on science (HOS) may afford students opportunities to better understand the nature of scientific research. The NRC (1996)7 suggests, “learning about science might improve students general understanding of science” (p 200). The use of historical perspective has been explored in other © XXXX American Chemical Society and Division of Chemical Education, Inc.
Received: July 16, 2017 Revised: January 9, 2018
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DOI: 10.1021/acs.jchemed.7b00524 J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Laboratory Experiment
is too sensitive to measure iron concentrations (0.5−5 mM) in modern supplements.
students make a connection between their laboratory experiments today and the history of chemistry research. Early physicians suggested that iron would cure a variety of diseases, considering iron to be a source of “vitality” and spa waters to supply the iron (Table 1). Today iron supplements
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EXPERIMENT OVERVIEW The goals of this experiment were to (1) see how natural materials like gall nuts can be directly used in a chemistry lab, (2) compare visual and digital forms of data to understand how physician chemists performed research before the age of quantitative analysis, (3) use volumetric techniques properly to make dilutions, (4) identify and visualize the relationship between absorbance and solution concentration through Beer’s law, (5) learn how to construct a calibration curve for interpolation, and (6) illustrate the importance of graphing as a form of scientific observation and data analysis. The experiment was implemented for four years (2013− 2016) in a first semester, general chemistry lab course for majors at Loyola University Chicago as well as a one semester course (2015) in a community college general chemistry laboratory. The first two years (2013−2014) were pilots of the experiment. Data from those years are not used in the student results below. This experiment could also be used in a second semester general chemistry class or an upper level analytical course. The experiment was written to follow a guided inquiry format in lab, specifically using the science writing heuristic template (SWH).26 Students read the prelab activity, the 1684 text of Short Memoirs for the Natural Experimental History of Mineral Waters of Sir Robert Boyle, and the procedure (prelab reading and student experimental procedure) to come up with a beginning question that could be tested and answered with an experiment. The procedure provided a general outline of the experiment, and allowed students to develop their own parts of the procedure in order to collect data to answer their beginning question(s). Students generally worked in pairs on the experiment, but they agreed with their lab group on procedure and data collection decisions. Students combined their data as a class to ensure replication and comparison of experimental results. Each student wrote a lab report that followed the SWH format.26 All reports were graded using the SWH rubric.26 The experiment was conducted during a regular 3 h lab period. Volumetric glassware, spectrophotometric techniques, and calibration curve concepts were introduced before the experiment as part of the prelab introduction
Table 1. Comparison of Iron Concentrations in Various Natural Sources of Water Location of Water Source
Iron Concentration mg/L
μM
Bath, Englanda Iranb
0.009−1.45
0.2−26.0
0.00212−0.621
0.0380−11.1
Turkeyc Chinad East Russiae Englandf
0.04−0.2 40.27 1.5
0.7−4 724.4g 27
0.011−1.5
0.2−27
Comments
Back calculation of national averages Well water Springs, cold water Razdolnov Spa Literature compilation of 40 water samples: Scarborough (1889) to Harrogate (1950)
a
See ref 18. bSee ref 19. cSee ref 20. dSee ref 21. eSee ref 22. fSee ref 23. gSimple dilutions could bring this sample into the color development range of tannic acid (0−200 μM).
are taken on the advice of modern day physicians. Those supplements can range from 0.5 to 5 mM iron in concentration. One supplement from Trace Minerals Research7 touts that the iron ion is the best form of iron for the body. This amount is much higher than those found in natural sources of water in Table 1. Little was known about spa waters and their efficacy as therapy in early times so early physicians began a set of protocols to determine what was in spa waters17 and how the properties of one spa differed from another (see also Table 1). Rudimentary tests of spa waters began ca. 1300, when Italian physicians evaluated precipitates obtained by evaporating spa waters (Hamlin, p 23).17 A common test for iron by color change of a gall nut extract was known in antiquity, first described by Pliny (23−79 C.E.).16,24,25 A gall nut is a small round ball formed (picture available in instructor directions in Supporting Information), typically on an oak tree, around a wound. It is high in tannic acid that can be extracted from the crushed or powdered gall. Use of the ground gall re-emerged from obscurity with the advent of spa water testing described by Paracelsus (The “Father of Pharmacology”) in 1562. Gabriel Fallopius (1523−1563, Paduan physician, for whom we derive the name Fallopian tubes) is credited24 with giving increased detail on the performance of the test. The first English description of the reagent occurs in a 1576 translation of the work of physician Conrad Gesner (1516−1565) published in England as The Newe Jewell of Health. Even the chemist, Sir Robert Boyle, renowned today for establishing a series of gas laws, was interested in the issue of spa waters. He is credited with converting the previously qualitative method into the first quantitative analytical chemistry method. This is the method that students will be performing in this laboratory experiment described here. When reacted with iron(II) ions within a specific concentration range and pH, the gall nut produces a deep purple color (Figure 4). The color development range (0−200 μM) occurs over the natural range of iron within domestic and mineral waters as shown in Table 1. Ironically the measurement
Calibration Standard Preparation
For the preparation of iron calibration standards, the students massed an initial amount of iron(II) sulfate heptahydrate (50.0 mg) and mixed this with 0.500 L of water. This amount provided students with a concentration of around 360 μM which is more concentrated than the amount of iron generally found in drinking water [Table 1 (0−200 μM)]. This allowed the standards to be made up within the linear range 10−40 μM. The preparation of samples was different depending on the teaching assistant in charge of the laboratory seen in Table 2. Even though preparation was different, students were able to obtain sufficient data for their calibration curves and interpolation. The gall nut solution was prepared the same day by students or the previous day by the teaching assistant (TA). It is not stable, and it cannot be prepared more than a day in advance B
DOI: 10.1021/acs.jchemed.7b00524 J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Laboratory Experiment
• The powder is also unstable and degrades with time. • Color intensity depends on the proportion of gall to iron. • The gall nuts produce their own color which makes monitoring the change difficult. That is, the background measurement must be accounted for. • While gall nuts can react with other metals, they do not react with arsenic. • The test can result in “false negatives” as some mixtures of metals containing iron will not react. Students arrived in lab with a beginning question and a tentative experimental procedure. Good beginning questions asked about the relationship between absorbance and concentration of iron in the solutions, although some students studied other variables, like the amount of gall nut solution used or the amount of time solutions are allowed to react. Some example beginning questions from students included the following: “How does the absorbance of the iron(II) sulfate heptahydrate−gall nut solution relate to the concentration of the solution? How does the concentration affect the magnitude of the absorbance?” “What is the relationship between concentration of iron and the absorbance of light in solution?” Ideally, the class decided on one question to answer with their experiment, although some groups answered different questions in addition to the main question. The answers to these questions, although simple, provided students with the opportunity to obtain data that could fully support their claim, the answer to their beginning question. For the procedure, students were prompted to determine the number of iron standards and how to make them. In addition, some students prepared the ground material and then capped it into a vial to prevent air oxidation. Some students ground several gall nuts together to create mixed powder for several groups to use so that can data might be cross-compared. Some students used litmus paper to test the pH. Although our students did not ask for other metals to test, they could check the specificity of the reaction (see Supporting Information on chemistry of the reaction). All of these are options in their procedure allowed for student driven inquiry.
Table 2. Students’ Preparation of Solutions Location and Year
Prepared Gall Nut Solution
Prepared FeSO4·7H2O Solution
LUC2013 LUC2014 LUC2015 HC2015 LUC2016
Yes Yes No Yes No
Yes Yes Yes No Yes
nor used again at a later lab period. Students prepared the gall nut solution by first pulverizing the gall nut with a mortar and pestle. To extract the tannic acid, the gall nut was placed in warm (not boiling) deionized water in a beaker. The gall nut rested for at least 30 min in the water bath. When the instructor prepared the solution, it was prepared the night before, and further instructions can be found in the instructor directions in the Supporting Information. All students brought a water sample from home to use as an unknown to interpolate from their calibration curve. The experiment allows for a range of options for students to explore; thus, there is variability in the students’ procedure. Students in our lab did not obtain the same results for a variety of reasons. Examples that explain this variation include the following: Students heated the gall nut solutions for different periods of time and may have extracted different amounts of tannic acid. They also added different amounts of indicator, the gall nut solution, to their iron standards based on what they perceived as dark, but still translucent enough to be detected by the spectrophotometer. Prompts in the lab procedure asked them to visually estimate the color to make sure it was not too dark for the spectrophotometer. Some students had to dilute their solutions if they found their initial solutions were too dark for the spectrophotometer. In 2014, the class chose a very wide range of calibration standards whereas students in 2015−2016 were encouraged by the TA/instructor to use a smaller range. Students also chose the highest peak (wavelength) on the spectrophotometer, which was not always consistent depending on the gall nut indicator. In the prelab, the students were asked to identify the maximum wavelength on a sample graph. Even though the procedure prompts students to choose a wavelength to determine the absorbance, this was the students’ first encounter with a spectrophotometer. They may not have chosen the highest absorbance each time even if the wavelength did not match the previous. Despite these variations, most students obtained linear calibrations curves, as discussed further below.
Materials Required
For several years, gall nuts were purchased from Griffin Dye Works;27 however, they are now available from various online retailers. In addition, tannic acid from Chinese gall nuts available through Sigma was used in 2016. Iron(II) sulfate heptahydrate or iron(III) nitrate was used from Sigma. Volumetric pipets, flasks, mortar and pestle, and any UV−vis spectrophotometer can be used. Individual Vernier spectrophotometers were used. Beakers and hot plates were required for preparing the gall nut solution.
Beginning Questions
Students were required to come up with a beginning question before arriving to class. In this respect, referring to Boyle’s comments from the 1600s was quite useful. The students were encouraged to read the text although the language was hard to interpret. It was to give students an opportunity to consider the difficulties experienced by early chemists/physicians in research. Student questions might come from annotations of Boyle’s comments: • The color of the gall nut iron mixture is pH dependent. • The test can be performed with gall nut extract, but the reaction will be slow and color varying. • Powdered galls work better but must be oxygen free. • The quality of the gall nut extract depends upon temperature. Cooler extracts work better. • The extract is unstable and degrades with time.
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HAZARDS Iron(II) sulfate heptahydrate or iron(III) nitrate is hazardous in case of skin contact (irritant), of eye contact (irritant), of ingestion, and of inhalation. Gall nuts are irritants and are not to be ingested. Students should be careful with hot glassware.
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RESULTS AND DISCUSSION: STUDENT RESULTS Table 3 shows how students performed the experiment as part of their laboratory class in three separate years. In addition to the laboratory class, students in a separate general chemistry course were asked in an extra credit assignment to try to replicate Boyle’s experiment at home (LUC2014 E.C). This is discussed after the class results. C
DOI: 10.1021/acs.jchemed.7b00524 J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Laboratory Experiment
Table 3. Comparison of Student Participation and the Number of Data Sets Obtained Semester LUC2014 E.C.c LUC 2014d LUC2015e HC2015f LUC2016e Total
Number of Data Points Collected
Number of Data Setsa
% Discarded Data Pointsb
7
NA
NA
NA
34 64 19 80 204
78 90 36 81 285
26 30 9 27 92
NAd 29 0 21
Students, N
a
Students per group ranged from 2 to 4. bData points were discarded if trending downward, flat, or beyond the linear range. cData from LUC2014 E.C. experiment were qualitative, and discussed below. d LUC2014 data were not used for analysis or in the figures below. e LUC students used 3 standards per data set. fHC students used 4 standards per data set. Figure 2. Experimental calibration curves for iron from an aggregate of student data from HC2015 black ● with error bars. HC data were consistent and was combined and can be shown with error bars because all groups agreed to use the same iron concentrations. LUC2016 (○ and ---) students as a class did not agree on iron concentrations. The red ● represents data from the literature.28
Each class added their data to a class database in Microsoft Excel. Data from 2015−2016 lab classes were compiled into calibration curves (Figures 1 and 2). Data from 2014 are not
separate groupings of data, one showing a more sensitive calibration curve (labeled “-a”) and one showing a less sensitive calibration curve (labeled “-b”). The less sensitive curve was the more accurate one compared to literature (the red data points in Figure 1). The variation in curves was not a function of the selected wavelengths (which ranged from 396 to 591 nm with 90% of the measurements >570 nm). The data suggest uniform error in calculations. Students were either not yet adept in calculating their dilutions and reported their dilutions incorrectly, or groups may have reported their data in grams of iron per milliliter instead of molarity. Another issue was students used the molar mass of iron sulfate, rather than the heptahydrate, to calculate concentrations. Because of these issues, adjustments to lab were made between LUC2015 and LUC2016. The TA introduced students to molarity calculations before and after performing the experiment. The TA also encouraged students to pay close attention to the name, formula, and molar mass of the chemicals they used in lab. The LUC2016 data are more consistent with each other and with literature(Figure 2). Figure 2 shows aggregate data from student data sets in LUC2016 and HC2015. Both sets of data fall nicely along the reported literature calibration curve. The range of linearity was found to be
