Article pubs.acs.org/jchemeduc
Do-It-Yourself Experiments for the Instructional Laboratory Norman C. Craig* and Cortland S. Hill Department of Chemistry and Biochemistry, Oberlin College, Oberlin, Ohio 44074, United States S Supporting Information *
ABSTRACT: A new design for experiments in the general chemistry laboratory incorporates a “do-it-yourself” component for students. In this design, students perform proven experiments to gain experience with techniques for about two-thirds of a laboratory session and then spend the last part in the do-it-yourself component, applying the techniques to an experiment of their own design. An emphasis on classifying inorganic reactions as acid−base, redox, complexation, and precipitation supports this program. An example is an enthalpy of reaction experiment in which students study an acid−base reaction, a complexation reaction, a precipitation reaction, and a redox reaction. Students perform three of the proven reactions and then do the fourth type with a reaction of their choosing. Other examples of experiments are described. The do-it-yourself component engages students in design and interpretation within a reasonable framework. With this approach, students take chances with new ideas as do working scientists. KEYWORDS: First-Year Undergraduate/General, Curriculum, Laboratory Instruction, Physical Chemistry, Inquiry-Based/Discovery Learning, Acids/Bases, Equilibrium, Oxidation/Reduction, Precipitation/Solubility, Qualitative Analysis
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been adequately trained. The general idea is that students do proven experiments for about two-thirds of the laboratory session and then work on an experiment of their own design, the do-it-yourself component, using the same equipment and methodology. Discussion of the do-it-yourself component is included in a lecture session prior to the pertinent laboratory. In addition, each of these experiments is preceded with an open book quiz or a prelab worksheet to ensure that students are prepared and aware of common pitfalls. An overall framework is useful to organize students thinking. In this work, the framework has been a sustained emphasis on four types of chemical reactivity: acid−base (Brønsted−Lowry) reactions, redox reactions, complexation reactions, and precipitation reactions in aqueous chemistry. In this context, net ionic reactions are emphasized and the idea that the parts of a substance, such as barium ions and chloride ions in barium chloride, may have different chemistries in these reaction categories is highlighted.
tephen Hawkes questioned the value of laboratory experience for students in lower-division laboratory programs.1 He regarded computer-based projects (or virtual laboratories) as better instruction because they emphasized design and decision-making. We made a similar argument for limited use of computer experiments long ago when computer applications were not widespread.2 Work on the computer is, however, no substitute for direct experience with actual materials. Who wants to be operated on by a surgeon who has only done exercises on a computer? Others have advocated and produced discovery or workshop laboratories in which students do research-like projects. In these laboratories, students assume considerable responsibility for design and interpretation, and they encounter failures, although hopefully not so many that they lose confidence and interest in doing chemistry. Such laboratory programs are laudatory but are hard to mount on a large scale with relatively inexperienced teaching assistants, as is true in the majority of university and college laboratories. In addition, even under favorable institutional circumstances, it is hard to find fresh ideas for discovery-based laboratories year after year and to sustain the program after the initial enthusiasm has worn off and the instructional personnel have changed. A discoverybased laboratory program may also suffer from not giving students sufficient exposure to the range of techniques and observational experience students should have. Is there a middle ground in which students learn needed techniques in an efficient way, do some design, and take some chances? In so doing, do they achieve greater ownership of the experiment and see the learning of techniques having an immediate payoff? We have had good experience with a “do-it-yourself” design of some experiments for the general chemistry laboratory that has students take ownership of the experiment by doing some design after gaining experience with techniques and a variety of chemical substances. This experimental design is adaptable to large laboratory programs, provided the instructional staff has © 2012 American Chemical Society and Division of Chemical Education, Inc.
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EXPERIMENTS
Enthalpies of Reaction
A concrete example of an experiment with a do-it-yourself component done in the context of instruction in thermodynamics is described. Students measure the enthalpy of reaction for examples of the four types of reactivity. They compare their experimental values for enthalpies of reaction with values computed from standard tables of thermodynamic data. Students chose to do three of the four proven experiments and then investigate an example of their own choosing and design for the fourth reaction type. The experiment is done on a small scale in nested Styrafoam cups with sensitive digital thermometry. Instructions are provided for an acid−base reaction (hydrochloric acid and Published: March 23, 2012 755
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Table 1. Classification of Reactivities of Chemical Species Precipitating Agent Speciesa Al(s)* Cu(s)* Fe(s)* Ni(s)* Sn(s)* Zn(s) Ag+ Ba2+ Ca2+* Cu2+ Fe2+* H3O+ Na+ Ni2+* Zn2+* Br− Cl− CO32−* HCO3− I−* NO3− OH− SO42− Br2(aq) H2O NH3(aq)
Acidb
Baseb
Oxidizing Agent (in V)b
Reducing Agent (in V)b
Cationicb
Anionicb
Complex Former Cationb
Ligandb
++ (0.76)
++
+
++ (0.80)
++ ++
++
+ (0.153)
++
++
+ (0.00)
+ (− 1.36) + (− 1.07)
+ +
+ +
+ (− 0.40)
++ +
++
+ ++ (0.96) ++ +
+
+ ++
+ (0.17)c ++ (1.07)
+ (− 0.45)d + (0.00) + (Source of OH−)
++ +
* designates species for which entries are to be filled in. bSymbol ++ means high degree of reactivity; symbol + means moderate degree; blank means little or no reactivity. cAs an oxidizing agent, reacts very slowly in acidic solution and even more slowly in basic solution; not a recommended reagent. dIn basic solution; less reactive in acidic solution.
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change. If the temperature change is too large, excessive thermal exchanges with the surroundings compromise the measurement. Other students approach the design of their part of the experiment from the theoretical side. They develop some ideas before the laboratory session by computing the ΔrH of a proposed reaction from tabulated data and then adjust their reagent concentrations to achieve an optimal temperature change of approximately 3 °C. Full instructions for the enthalpy of reaction experiment are in the Supporting Information. Well-stocked reagent shelves of interesting solid chemicals and solutions of known concentrations need to be available for an experiment such as the enthalpy of reaction experiment. Doing the reactions on a small scale makes supplying the varied reagents manageable and reduces concerns about students doing experiments with unfamiliar chemistry. Good quality tables of thermodynamic data must be on hand. The instructional staff must be capable of helping students interpret surprising results for a wide range of reactions and of helping students find new possibilities when they are disappointed. In all the laboratory sessions, faculty members are present as are upper-class teaching assistants, who have personal experience with do-it-yourself experiments and who have attended a training session for the specific experiment.
solid sodium bicarbonate), a precipitation reaction (aqueous copper nitrate and aqueous sodium hydroxide), a complexation reaction (aqueous zinc nitrate and aqueous ammonia, an intentionally colorless reaction), and a redox reaction (solid ferrous sulfate with aqueous potassium permanganate and sulfuric acid). The use of the reaction of solid sodium bicarbonate and aqueous sodium hydroxide was a student contribution in response to a request for suggestions of an endothermic reaction to counter the notion that only exothermic reactions are spontaneous. The digital thermometer, which depends on a rugged diode sensor and reads to 0.01 °C, is of local design. Solutions of appropriate concentrations are provided. Students are asked to choose and do three of the available experiments. In each case, they must supply a qualitative test to show which reactant species was the limiting reagent for use in computing the enthalpy change per mole of reaction. For the fourth type of reactivity, the one they chose to omit from the proven set, students design and perform an experiment involving different chemistry than is described in the instructions. They know from discussion in class of two ways to approach this part of the project. One approach is to try out their idea in the laboratory. In doing so they may be surprised to find an almost thermoneutral reaction such as the one between aqueous acetic acid and aqueous sodium hydroxide. Of course, a ΔrH of zero is a valid possibility for this quantity, which spans the range from negative to positive values. For reactions with significant positive or negative values of ΔrH, it is important to have a reasonable temperature
Reactivity of Species
Another, more qualitative experiment with an emphasis on design and testing ideas helps prepare students for the enthalpy of reaction experiment. This experiment is built directly on the theme of the four types of reactivities. Students have learned a 756
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a convenient equimolar mixture of quinone (ox) and hydroquinone (red), and a platinum electrode give convenient sensitivity to hydrogen ion concentration in acid solutions. Students use the Nernst equation and the measured emf’s (electromagnetic force) to find the concentrations of an electrode-active species. From this tiny concentration, the excess concentration of the reaction partner, and the stoichiometric concentration of the product, students compute the equilibrium constant. In the do-itself-yourself component of this experiment, students omit the supplied experiment for one type of reactivity and design a measurement of the equilibrium constant for a reaction of the same type. The use of the mini-cell design enables this experiment. The needed quantities of a range of reagents are supplied in dropper bottles. Students are free to work with a wide range of systems, including metal ions, because the quantities are so small that hazards are at a minimum. Results are compared with entries in the tables of the Handbook of Chemistry and Physics.4 Measurements with the mini-cells, which have a short distance between electrodes and low resistance, give equilibrium constants within an order of magnitude of the accepted values. Students are impressed with the wide dynamic range of electrochemical cells, which permit the measurement of tiny concentrations with simple equipment without disturbing the equilibrium. The chemical equilibrium occurs in one-half cell (one eyedropper) whereas the overall electrochemical reaction is held in electrochemical equilibrium.5 Full instructions for the galvanic cells experiment is in the Supporting Information.
hierarchy for the four types of reactivity. Redox reactions, which are often accompanied by acid−base chemistry or the other two reactivities, are at the top of the hierarchy. Complexation reactions are second, followed by precipitation reactions. Acid− base reactions are at the bottom. Acid−base chemistry may accompany any of the other three categories. Students are supplied with a table in which the reactivities of individual species, mostly ions, are classified as acid, base, reducer, oxidizer, complexer, ligand, precipitating cation, or precipitating anion. Table 1 shows that the reactivities of the species are supplied in qualitative terms with double pluses, single pluses, and blanks. For acid−base, redox, precipitation, and complexation chemistry, semiquantitative guidance comes from tables of pKa’s, E°’s, Ksp’s, and Kform’s, respectively. For about 60% of the species in the table, estimates of their reactivities are supplied. For the remaining species, students complete the table before attending the laboratory session. In categorizing these species, students gain valuable practice with important tables of data. As Table 1 implies, a number of species have multiple reactivities that depend on the nature of the reaction partner. Students also confront how acidity impacts spontaneity of the redox chemistry of oxygenated species such as the nitrate ion. In the laboratory session, students use species from anywhere in the table to do reactions on a small (10 mm × 70 mm) test tube scale for examples of each of the four reaction types. An immediate challenge for work with most species is to choose a counterion that is inert in the reaction of interest. Of course, inert cations are easier to find than inert anions. Students must supply convincing visual evidence that the proposed reaction occurred. Changes of phase and color help. Supplementary reactions, such as pH indicator paper or the extraction of iodine or bromine into a hexane layer, may be used in the demonstration of the reactivity. Many surprises occur for the students and for the laboratory staff in these experiments. Of course, an expected type of reaction, such as a supposed complexation reaction between copper(II) ion and iodide ion, may be something quite different. In this case, cuprous iodide precipitates and triiodide ion forms. Consultation with the laboratory staff and among the laboratory staff is frequent. As for the enthalpy of reaction experiment, well-stocked shelves of solids and solutions are needed. The small scale of reactions keeps the demand for materials low, which are mostly available in dropping bottles, and minimizes concerns about hazardous materials and reactions. Full instructions for the classification of species experiment are in the Supporting Information.
Semi-Micro Qualitative Analysis
Semi-micro qualitative inorganic analysis is another vehicle for putting students on their own to select and design experiments. The do-it-yourself outcome in this familiar experiment is achieved by providing students with a full scheme for analyzing 26 different cations and 10 different anions but then supplying unknowns that are expressly limited to a small number of possible designated species. Often only a single cation or anion from a group or subgroup is a possibility in a given unknown. Unknowns that are limited in this way can be done in less than half an hour by a student who plans ahead and adapts the standard methods to the limitations of a particular unknown. In contrast, slavish adherence to the stepwise procedure can take hours. Of course, the reactions and equilibria used in inorganic qualitative analysis are splendid exercises in the four types of reactivities emphasized in the three experiments already described. While the emphasis is on adapting known procedures to the efficient identification of unknowns, the whole program provides students with a memorable exposure to standard inorganic chemistry of aqueous solutions. As a summarizing experiment, students receive a hot or cold pack substance (a colorless binary salt, equivalent to a commercial product) for exploration of the thermal effect with added water and for identification of the ingredient. In designing their analysis method, students consider the solid’s lack of color and recognize that a commercial product must be inexpensive and nontoxic. Finally, they measure the enthalpy of reaction with water as a confirmation of the substance.
Use of Galvanic Cells to Measure Equilibrium Constants
A galvanic cells experiment contains full instructions for measuring equilibrium constants for three of the reaction types: a precipitation reaction, a complexation reaction, and an acid−base reaction. Because electrochemical cells depend on redox chemistry, it was redundant to include the measurement of an equilibrium constant for a redox reaction. These measurements are made in mini-cells constructed from glass eyedropper tubes inserted into a salt bridge solution of ammonium nitrate in a small vial. This mini-ware for galvanic cells has been described in the Journal.3 The prescribed precipitation equilibrium is the solubility product for silver chloride. The complexation equilibrium is the formation of the tetraamminezinc ion complex. The acid−base equilibrium is for the second dissociation of sulfuric acid. Quinhydrone, which is
Context
In the recent laboratories, the three experiments first described are done by pairs of students. The conversations about the design of the experiment are of high value. Joint responsibility 757
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for the do-it-yourself part blunts any tendency for students to adopt a parallel processing mode to hurry through the set part of the experiment. For the qualitative analysis, the students work individually. The do-it-yourself design has transferred successfully as the faculty members and student laboratory assistants have changed. This design works as well for a one-semester, advanced placement course or for the two-semester general chemistry course, in which most students enroll.6 In student evaluations of the laboratory program, the experiments with the do-it-yourself design receive more praise in comparison with other experiments without this feature. Thus, students recognize the value of the do-it-yourself feature and become more engaged by it.
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ASSOCIATED CONTENT
S Supporting Information *
Instructions for the experiments described in this article, other than for the qualitative analysis and the hot and cold packs experiment. This material is available via the Internet at http:// pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS Many Oberlin chemistry colleagues and students have participated in developing these experiments and working with them in the laboratory. Terry Carlton played a major role in developing the Reactivity of Species experiment. William Renfrow introduced the mini-cells made from eyedroppers. William Mohler designed and built the digital thermometers. Andrew Hare, a student at the time, suggested the endothermic reaction of solid sodium bicarbonate and hydrochloric acid.
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REFERENCES
(1) Hawkes, S. J. Chem. Educ. 2004, 81, 1257. (2) Craig, N. C.; Sherertz, D. D.; Carlton, T. S.; Ackermann, M. N. J. Chem. Educ. 1971, 48, 310−313. (3) Craig, N. C.; Ackermann, M. N.; Renfrow, W. B. J. Chem. Educ. 1989, 66, 85−86. (4) Handbook of Chemistry and Physics; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 2007. (5) Thompson., R. Q.; Craig, N. C. J. Chem. Educ. 2001, 78, 928− 934. (6) Further information about the laboratory program can be obtained from Cortland Hill at
[email protected].
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