In the Laboratory
Constructing Chemical Concepts through a Study of Metals and Metal Ions Guided Inquiry Experiments for General Chemistry Ram Lamba* and Shiva Sharma Department of Chemistry and Physics, Inter American University of Puerto Rico, P.O. Box 191293, San Juan, PR 00919-1293 Baird W. Lloyd Center for Chemical Education, Miami University Middletown, Middletown, OH 45042
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Each experiment introduces and associates several closely related chemical concepts.
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Students engage in decision-making processes in designing and carrying out procedures.
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Students predict possible outcomes of experimental measurements before performing the actual measurements.
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Students work cooperatively in small groups to collect data. The groups collaborate with one another to create a class data set for final analysis.
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Students use the data of the experiments to test their perceptions and to develop meaning of the concepts.
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Teachers lead students in discussion during the progress of the experiment that asks them to think about the implications of the observations and measurements being made.
The five experiments described in this paper may be done sequentially in successive weeks, or may be spaced throughout the semester or year of study. The inquiry nature of these experiments, including the intensive instructor–student and student–student interaction and discussion, allows these experiments to be done before the concepts are discussed in class. Experiment # 1: Are All Pennies the Same?
Procedure Students begin by seeking observable differences among pennies in a set that might help them answer the question, “Are all pennies the same?” When the class data are pooled, students find that a significant difference in mass exists between pennies minted before 1983 and those from that year onwards. A graph of mass as a function of year minted using the combined class data confirms this conclusion. Students then measure the total mass and total volume of sets of 3, 6, 9, or 12 pennies that they believe are similar in properties. When they pool their data to generate a single graph of mass versus volume of each set, they find that the two lines have different slopes (Fig. 1). Once again the data separate into two groups, based upon the date relative to 1983. The instructor suggests that students use a file to remove some metal at four or five points on the edge of a penny from each group, then cautiously slide the pennies into a beaker with about 60 mL of 3M HCl, cover the beaker, and set it aside. Almost immediately small bubbles will be observed on the edge of one of the pennies. If left undisturbed until the next period, the newer penny will be floating, whereas the pre-1983 penny remains on the bottom of the beaker. If the instructor wishes to demonstrate the floating penny effect within a single lab period, a post-1982 penny filed at 5 to 6 points floats in about 2 hours in approximately 60 mL of a 6M HCl solution.
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This experiment is an adaptation and expansion of one first developed at Holy Cross College (3). It introduces students to the differences between extensive and intensive *Corresponding author.
quantities and the usefulness of intensive properties for distinguishing between different substances by studying the intensive and extensive properties of zinc and copper.
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As appreciation for the value of emphasizing inquiry in general chemistry courses grows, there is a pressing need for experiments that employ these methods. While inquiry experiments come in a variety of formats (1, 2), all involve students firsthand in the scientific process by having them collect data from which they extract meaning. This paper describes a set of experiments designed to help students develop an understanding of basic chemical concepts within the framework of studying the properties and reactivity of metals and metal ions. The content of these experiments may be familiar to instructors. Unlike more traditional experiments, however, the instructional objectives of these experiments emphasize the construction of meaning from observation, measurement, and data analysis. The instructional format involves intensive student–instructor and student–student interaction and discussion. Each experiment in this series presents a question that students answer by performing experiments and making measurements. The experiments in the set have several common features:
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Figure 1. Are all pennies the same? Typical data.
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In the Laboratory 300 250
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Data Analysis During the postlab discussion students reflect upon the differences in properties of the pennies, the difference between intensive and extensive properties with respect to their use in distinguishing different substances, and behavior versus composition of the pennies. Students use their own data to justify their responses to questions about the nature of these differences. They also have the opportunity to learn the advantages of using large sample sizes to show trends in relationships of variables.
Mass Zn = Molar mass Zn Mass Mg Molar mass Mg The implications of this relationship is probed by rearranging this equation to the form
Mass Mg Mass Zn = Molar mass Zn Molar mass Mg Different results are obtained by comparing the ratio by mass of magnesium and aluminum needed to prepare a given volume of gas. In this case, the ratio by mass of the two metals differs from the ratio of their molar mass by a factor of 3/2,
Mass Mg 3 Molar mass Mg = × Mass Al 2 Molar mass Al
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Figure 2. How much is enough? Typical data.
the graduated cylinders, sample sizes may be decreased and the amount of hydrogen gas significantly reduced. However, the experiment has been run successfully and safely a number of times with the low-cost graduated cylinders. Fifty milliliters of HCl is poured into a 125-mL filter flask equipped with a rubber stopper. A piece of flexible tubing is connected to the sidearm. The graduated cylinder is filled with water and the top is sealed with Parafilm, making sure no air is trapped in the cylinder. A large container is filled about half full of water. While holding the Parafilm with one hand, the cylinder is inverted and lowered into the large container. When the top of the cylinder is below the level of water in the container, the Parafilm is removed. The open end of the flexible tubing is inserted into the cylinder and the cylinder is clamped into position. At this point the metal is quickly added to the acid and the stopper immediately placed over the flask. When magnesium is being used, since the reaction is extremely fast, it is recommended that the flask be inclined carefully and in such a way that the metal can be placed on the neck of the flask, providing ample time to place the stopper before the metal comes into contact with the acid. The instructor should advise the students on the appropriate methods of chemical disposal.
Data Manipulation and Analysis Once all data for the set of reactions has been collected, each student prepares a graph that plots the volume of gas as a function of the mass of metal consumed. The data for all of the metals are combined on the same graph (Fig. 2). The graph is used to find answers to the following questions: •
What happens to the volume of gas produced in this reaction as the mass of the metal sample increases? Is the same trend observed with all three metals?
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If you start with the same amount of different metals, do you get the same amount of gas?
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What mass of each metal is needed to produce 50 mL of gas? 75 mL? 100 mL?
which translates to a 1.5 to 1 mole ratio of the metals needed to prepare a given volume of hydrogen gas.
Procedure One group of students is assigned to study the reaction of metallic zinc with hydrochloric acid while others use either magnesium or aluminum. Each group works with 6– 8 samples of different mass. An excess of acid is always used with masses of metal that range from 0.02 to 0.2 g. Fifty milliliters of 1, 2, and 4M HCl could be used with Mg ribbon, Al foil, and Zn granular (20 mesh), respectively. The hydrogen gas is collected by water displacement using a large graduated cylinder. The actual sample sizes for the metal are adjusted according to the size of cylinder available (0.20 g of zinc yields approximately 80 mL of gas; a similar sample of magnesium and aluminum produces near 200 and 285 mL, respectively). The amount of hydrogen gas produced by these samples is sufficient to require very strict adherence to safety rules. If gas burets are used in place of
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Experiment # 2: How Much Is Enough? Students examine the quantitative aspects of the reaction between metals and hydrogen ions to experimentally develop and understand the concept of the mole and stoichiometric calculations that can be made using this concept. This knowledge will be reinforced and extended in experiment 3—How Much Is Too Much?—as students develop the concepts of limiting reagent, excess reagent, and spectator ion. Students explore the stoichiometric relationship between the mass of zinc, magnesium, or aluminum consumed in a reaction with a strong acid and the volume of hydrogen gas produced. Data for the class are combined to construct a single graph that enables the students to predict the mass of each metal required to prepare a given volume of gas. By comparing the results obtained with zinc and magnesium, the students find that the ratio by mass of the samples of these metals required to prepare a given volume of gas is virtually the same as the ratio of the molar masses of these elements.
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Students use the information obtained from the answer to the last question to find the answers to three more questions: •
What happens to the graph when the data for the 50-, 75-, and 100-mL volumes are plotted as a function of the ratio of the mass to the molar mass of the metal instead of as a function of the mass of the metal sample?
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What is the relationship between the average ratio of the mass of magnesium and zinc required to prepare the same amount of hydrogen gas vs. the ratio of the molar mass of the metals?
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In the Laboratory •
What is the relationship between the average Mg/Al mass ratio and the ratio of the molar mass of these elements?
To find the answers to these questions, students perform three more sets of calculations. First, they divide the sample masses by the molar mass of the metal. These values for all of the metals are plotted against the corresponding volume of gas on a single graph. Students find that the data for Mg and Zn fall virtually on the same line, but those for Al fall on a different line. Instructors lead students to realize that the calculations have converted the masses to moles of metal. From this, students can reason that the stoichiometry for the reactions of Zn or Mg with acid is the same but because the Al data fall on a different line, the stoichiometry of that reaction must be different from the other two. The masses and mass ratios of the varying elements (Zn/Mg, Mg/Al) needed to produce given volumes (50, 75, and 100 mL) of hydrogen are determined. From these results students can conclude that the stoichiometry of these reactions is similar to that of the earlier ones. They are then asked to write balanced net ionic equations to represent the reactions of each of the metals with acid, using the mole ratios that they have calculated.
Extensions This experiment can be connected to the chemistry of metals observed in everyday life. It can be used, for example, to explain why it isn’t a good idea to steam-clean “mag” wheels on a car (because magnesium is sufficiently active to react with the extremely small concentration of hydrogen ions in steam) or to explain why aluminum pots and pans look cleaner after having been used to cook acidic foods like sauerbraten (the “sauer”, or acid, in this food is strong enough to remove some of the metal from the surface).
magnesium with HCl and magnesium with H2SO4, students are asked about the role of Cl { and SO4 2{ as compared to H+ in the reactions.
Procedure The same apparatus is used for gas collection as was used in the previous experiment. This gives students the opportunity to refine their skills with a particular set of equipment and a particular procedure. In addition, they are able to concentrate more fully on the chemistry of the process because the procedure is sufficiently familiar for them to feel confident about making the necessary measurements. Each pair of students uses one metal and one acid and collects at least four pieces of data. Class data are pooled for analysis. Some possible combinations of reagents include 1. 0.05-g samples of magnesium ribbon with 5 mL of 0.1, 0.5, 0.7, and 0.9 M HCl. 2. 0.05-g samples of magnesium ribbon with 5 mL of 0.3, 0.6, 0.8 and 1.0 M HCl. 3. 0.05-g samples of magnesium ribbon with 5 mL of 0.1, 0.3, 0.5 and 0.7 M H 2SO 4. 4. 0.05-g samples of magnesium ribbon with 5 mL of 0.2, 0.4, 0.6 and 0.8 M H 2SO 4. 5. 5.0 mL of 4 M HCl with samples of aluminum foil having masses of approximately 0.04, 0.12, 0.2, and 0.28 g (weighed to the nearest 0.01 g). 6. 5.0 mL of 4 M acid with sample masses of 0.08, 0.16, 0.24, and 0.32 g aluminum foil.
Combinations of 1 and 2, 3 and 4, or 5 and 6 provide a set of eight data points for a metal, the optimum amount for analysis. The instructor might prefer to use only one of the combinations. Either arrangement will serve the purpose of demonstrating the concept of limiting reagent. How-
Experiment # 3: How Much Is Too Much?
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recognize that the observable characteristics of a chemical reaction involve a limiting reagent
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recognize, identify, and distinguish between limiting reagents and excess reagents
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recognize experimental conditions under which the identity of the limiting reagent can change
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recognize and distinguish between spectator ions and reactant ions
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Two types of study are made, using the reaction of magnesium with hydrochloric acid and sulfuric acid and the reaction of aluminum with hydrochloric acid. The purpose of the former is to determine the effect of varying acid concentration on production of hydrogen gas by magnesium. In the latter study, students investigate the effect of varying the mass of the aluminum metal on the production of gas with 4M HCl. Students are asked to create graphs to predict outcomes of these studies. (They often draw one straight line for the volume of gas as a function of metal mass and fixed amount of acid, which suggests the volume will increase continuously.) They are also asked to predict what happens when the concentration of HCl added to a fixed amount of metal is increased. (Some predict that the volume of gas increases as concentration increases; others predict that the change in concentration will have no effect on the volume of gas liberated.) In comparing the graphs of
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This experiment introduces the concept of limiting reagent and refines and further develops the concept of stoichiometry developed in experiment 2 by helping students learn to:
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Figure 3. How much is too much? A: 5 mL of 4M HCl with Al foil. B: 0.05 g of Mg with 5 mL of acid.
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In the Laboratory ever, combinations 1 and 2 together with 3 and 4 must be used by students in order to develop the concept of spectator ion.
Data Manipulation and Analysis When all the class data are collected, students prepare a graph showing the volume of gas collected versus the concentration of the acid for magnesium and/or one showing the volume of gas collected as a function of the mass of aluminum used, depending upon the approach(es) chosen for study. Sample of such graphs are shown in Figure 3. The discussion of this experiment can be guided toward introducing the concept “limiting reagent” by questions such as the following: •
What happens to the volume of hydrogen gas collected as the mass of metal is increased?
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What happens as we increase the concentration of acid added to a fixed mass of metal?
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Will these trends continue indefinitely? What is your evidence?
Students are asked to reflect on any differences between their predictions and their observations and to explain these differences, if necessary. During the discussion, students use their data to conclude that the amount of one reagent in a chemical reaction can limit the amount of product formed. Furthermore, they discover that the identity of the limiting reagent changes as the mass of metal or the concentration of acid increases. Understanding of the characteristics of a limiting reagent can be enhanced and reinforced by using questions such as the following: •
Which is the limiting reagent when the mass of metal or the concentration of the acid is relatively small?
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What reagent is present in excess when we add only a small sample of the metal or a relatively dilute solution of acid?
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Does the identity of the limiting reagent remain the same throughout the experiment? Why or why not?
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What can we look for when we encounter a problem, to decide whether the problem involves a limiting reagent?
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What should we look for to identify a limiting reagent?
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Is there any chemical reaction that can be run with no single limiting reagent?
Experiment # 4: What Is an “Active” Metal? For years, chemists have organized metals and their ions in tables of relative activity. This experiment introduces students to the observations, knowledge, and logic used to sort common metals into an activity series. Students learn that some substances are more active than others in the sense of being more reactive in a chemical reaction. They learn in experiment 1 that zinc metal reacts with hydrochloric acid, whereas copper metal does not. Zinc, therefore, can be labeled as more “active” than copper. In this experiment students learn to avoid the mistake of confusing metals and metal ions, take a first step toward developing a table of oxidation–reduction half-reactions, and gain the experience necessary to discuss oxidation–reduction reactions.
Procedure Students start by examining what happens when pieces of iron metal are immersed in solutions that contain either Fe2+ or Cu2+ ions. They then repeat the experiment with strips of copper metal. No changes are observed for three of the four combinations, but something seems to hap-
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pen to the iron strip when it is immersed in a solution of the Cu2+ ion. The instructor leads a discussion that helps students deduce from these observations that two things occur: as copper ions receive electrons to form copper metal atoms, the iron metal must be giving up electrons, thus becoming iron ions. Apparently, the reverse reaction does not occur. Students are asked, “How can these observations be used to decide which metal is the more active?” Students design and carry out experiments that allow them to compare the relative activity of zinc and magnesium with that of iron and copper. All they need to do is recognize that the more active of a pair of metals reacts with solutions that contain ions of the less active metal. They perform these experiments with freshly cleaned strips of metal and a set of 1.0M aqueous solutions, each containing one of the metal ions. In addition, they immerse a clean strip of each metal in a 1.0M solution of hydrochloric acid to observe which metals react with the acid and which do not.
Data Manipulation and Analysis Students are asked to list the four metals in order of decreasing activity and to give evidence supporting the placement of each metal in the series. They are asked to use their observations to place H2 in the series as well. The instructor can use questions such as those listed below to guide the students in these tasks. •
How did you decide whether one of the metals reacted with hydrochloric acid?
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Which is the least (or most) active metal in this experiment? What observation(s) convinced you that this is the least (or most) active metal?
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Where do the other metals fall in the activity series?
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Based on your observations, where would you place H2 in the activity series?
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What can we do with an activity series once it has been established?
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What general rule can be used to predict whether an electron transfer reaction is spontaneous?
Extensions The knowledge gained from this experiment can be used to discuss questions related to other metals, such as why aluminum soft drink cans and “tin” fruit juice cans have to be lined with plastic. Most students believe that aluminum metal is “inert” to chemical reactions because of their experience with products such as aluminum foil and aluminum siding. They might explore the reasons for the discrepancy between this belief and their experimental results. Finally, they might extend these ideas to explain the chemistry of batteries or the form in which metallic elements are found in rocks and minerals. Experiment # 5: How Active Are the “Active” Metals? In the final experiment of the series, students develop a quantitative model by linking cell voltage and metal activity, develop a deeper understanding of the concept of a spontaneous chemical reaction, and form the basis for discussing the relative strengths of various oxidizing and reducing agents.
Procedure Students prepare a voltaic cell using zinc and copper metals and 1.0M solutions of the metal ions. The connections between the metal strips and the voltmeter are arranged so that the voltage reading on the meter is positive.
Journal of Chemical Education • Vol. 74 No. 9 September 1997
In the Laboratory A discussion leads the students to the idea that the positive sign corresponds to a spontaneous chemical reaction and that the magnitude of the cell voltage gives information about the driving force behind this reaction. The larger the voltage for a cell, the larger the force driving the reaction towards the products of the reaction and the larger the difference between the activities of the two metals. Questions are used to make sure the students understand the meaning of the numbers they measure with the voltmeter: •
When zinc metal reacts with Cu2+ ions, which substance is oxidized and which is reduced?
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In which direction do the electrons flow in this cell, from the zinc metal to the copper metal strip, or vice versa?
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What would you conclude if you replaced the zinc strip with another metal and found that the cell voltage changes sign?
Once students show they understand the meaning of the numbers they read for the cell voltage, they are asked to design and perform a set of experiments to compare the activities of iron, magnesium, and zinc with copper metal. They report the sign and magnitude of the cell voltages and write balanced net ionic equations for each of the reactions studied.
Data Manipulation and Analysis Students are asked to list the reactions in order of increasing value of the cell voltage and the metals in order of decreasing activity. The following questions may be used to guide the discussion:
quantifying reaction-driving forces when they are first introduced to the interactions of energy and matter in chemical reactions. Conclusion The inquiry-based general chemistry curriculum developed at Inter American University of Puerto Rico has brought about changes in the role of both student and instructor. By collaborating with their peers to acquire and interpret data, the students inevitably accept more responsibility for their learning. The instructor is no longer the source of knowledge to be learned, but a guiding force in the learning process—choreographing the activity and its discussion, ensuring that student creativity is called for in small, manageable increments. The inquiry-based general chemistry curriculum inculcates student learning using easily accessible materials from daily life through direct experiences. Although our model is not completely open-ended, there are sufficient elements of creativity in this structured format to support the insights expected of students in the process of discovery. Preliminary evaluation of these experiments at both the Inter American University of Puerto Rico and Purdue University suggests that we have met our goal of transforming the traditional laboratory experience into one in which some of the burden for learning has been shifted from the instructor to the students. In our curriculum, students take an active role in the planning and carrying out of the activities, with emphasis on the processes, not the products, of science.
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What does the magnitude of the various cell voltages measured in this experiment tell us?
Acknowledgments
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What can we deduce from the fact that the magnitude of the voltage for the zinc/copper reaction is approximately 1 volt?
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Which cells give voltages that are larger than the zinc/ copper reaction? Which ones are smaller?
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What would happen to the magnitude of the cell voltage if we compared zinc with iron instead of copper?
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What would happen to the sign of each of the voltage measurements in this experiment if we compared each of the other metals to a sample of iron instead of copper?
We are grateful to the National Science Foundation grant DUE-9354432; the U.S. Department of EducationMSIP Program, grant P120A30018; and the Inter American University of Puerto Rico for partial funding of this work. We are also grateful to Nancy Konigsberg-Kerner, Lecturer and Coordinator, General Chemical Lab Program, Chemistry Department, The University of Michigan, Ann Arbor, Michigan, for a critique and suggestions for this publication.
Extensions The results of this experiment can be connected with a wealth of information about the behavior of metals, some of which has been described in previous experiments. While this experiment can be extended to fairly sophisticated ideas, it is also possible to use it reasonably early in a first-year course, because the primary purpose is to collect quantitative information to make more precise comparisons than are possible with qualitative data. Students will be able to make use of the data collected without knowing all the topics usually covered in electrochemistry. The benefit of doing this early is that students gain experience with a method of
Dedication The principal author wishes to dedicate this publication to Waldemar Adam, Institute of Organic Chemistry, University of Wuerzburg, Germany on his 60th birthday. Professor Adam was one of the forces that inspired me to go into chemical education. Literature Cited 1. New Directions for General Chemistry; Lloyd, B. W., Ed.; Division of Chemical Education of the American Chemical Society, 1994; pp 34–36. 2. Ditzler, M. A.; Ricci, R. W. J. Chem. Educ. 1994, 71, 685. 3. Ricci, R. W.; Ditzler, M. A. J. Chem. Educ. 1991, 68, 228.
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