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New Approaches to Process and Content in Introductory Chemistry
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by James N. Spencer
Franklin & Marshall College
Lancaster, Pennsylvania
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| | The traditional classroom setting has its origins in the medieval university. A master --the only one who had access to books -- lectured, students listened and recorded. At the beginning of the last century John Dewey pointed out that the problem with the traditional education was that it was all made for listening.1 One of the difficulties with teacher-talk is that students don't make the same connections as instructors. The belief that students will see the link between the instructor's words and the concept derives from the inability to conceive of a different point of view. Research has shown that telling is not teaching, that it is not possible to transfer an idea intact from the mind of the instructor to that of the student.2
Several pedagogical models have been developed to overcome these shortcomings of the traditional classroom experience. One such model is the guided inquiry instructional approach. The guided inquiry model uses the constructivist theory combined with the learning cycle in a group learning environment. The constructivist principle is that students construct their own knowledge based on what they already know. The learning cycle, consisting of three steps, mimics the scientific method. In the first step, students conduct an exploration. This step could involve a model or data, a demonstration, a case study or a variety of other hands-on activities. The second step, concept invention, is usually done in a discussion format. The third step is an application of the concept. The acquisition and application of knowledge has been shown to be a social act, thus the well documented group learning paradigm has been combined with constructivism and the learning cycle to produce a classroom pedagogy called guided inquiry.
The premise of the guided inquiry approach as developed at F&M3, 4 is that students will learn better when:
- They are actively engaged and thinking in class.
- They construct knowledge and draw conclusions themselves by analyzing data and discussing ideas.
- They learn how to work together to understand concepts and solve problems.
In the F&M implementation there are no lectures -- students work in groups of four with assigned roles, and the groups use guided inquiry worksheets to develop and learn concepts. The instructor serves as a facilitator and answers no question that the students could reasonably be expected to answer themselves. The worksheets provide students with data to be analyzed or understood. Critical thinking questions guide students to develop and understand concepts by asking questions scientists themselves would ask about the information or model supplied. The students then demonstrate understanding of the concept by applying their knowledge to solving problems. Examples of these worksheets may be found under "Downloadable Activities" on the Process Oriented Guided Inquiry Learning Web site, below.
It is not always immediately evident how an activity may be designed to lead students through the learning cycle to the development of the concept. In order to preserve the exploration step of the learning cycle it may be necessary to find non-traditional ways of presenting the concepts. One such example is electron configurations. In the traditional presentation of atomic structure, the first two steps of the learning cycle are reversed; that is, the concept, Schrödinger's model, is introduced and data are presented to support the model. In the guided inquiry approach, experimental data obtained from photoelectron spectroscopy are given and the student continually develops and refines models that lead to the concept. (Details of this approach may be found in references 5, 6, and 7, below.) Basically students use ionization energies for all the electrons in an atom to identify the various energy levels of an atom. The first ionization energy of helium is 2.37 MJ/mol, about twice that of hydrogen (1.31 MJ/mol). These energies are consistent with the higher nuclear charge of helium and therefore the electrons in hydrogen and helium are assigned to the first (n=1) energy level. Lithium has two ionization energies, 6.26 and 0.52 MJ/mol. Because of the increased nuclear charge of Li, it might be expected that the first ionization energy of Li would be greater than that of He. However, the lower first ionization suggests that one of the Li electrons is at a greater distance from the nucleus; that is, in the n=2 energy level. Boron has three ionization energies, Al has five, and K has six. These data allow the construction of the electron configurations of these atoms. Additional data complete the electron configurations through the transition metals.
The data used to develop electron configurations can also be used to define electronegativities. By using data already familiar to them (the ionization energies), students can relate the ability of an atom to attract electrons to itself in a molecule to the strength with which that atom holds onto its electrons. The bond type triangle, which shows a plot of the average electronegativity of the atoms in a binary compound versus the difference in electronegativity, allows the characterization of the bond type for a compound.6, 7, 8
Thus ionic, covalent, and metallic bonds may be discussed from a single diagram and the continuum that exists between these types of bonds made evident. The properties of compounds based on their position in the bond type triangle may be predicted. Students recognize the elements most likely to produce a ceramic, a semiconductor, or a brittle, high melting insulator. Ionic character increases from bottom to top of the triangle. For compounds with the same electronegativity difference, the bond type may be metallic, ionic, or covalent.
| Bond-type triangle for selected compounds. Such triangles may be used to classify solids as molecular, network covalent, ionic, or metallic. |
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Some years ago, Sanderson9 pointed out that the use of standard enthalpies of formation to calculate enthalpy changes for reactions was not particularly informative to students. This is due in part to the nature of enthalpies of formation, which refer to the enthalpy changes produced when the elements forming a compound are produced in their standard states of aggregation. Because of this convention, enthalpies of formation do not give direct information about the strengths of bonds in a compound. To use an example from Sanderson, the enthalpies of formation of CaCl2 and WCl2 are –796 and –255 kJ/mol, respectively. Which has the greater bond enthalpies? The bond dissociation enthalpies (the enthalpy required to separate the compounds into their gaseous atoms) for CaCl2 and WCl2 are 1217 and 1347 kJ/mol, respectively. Thus, because enthalpies of formation have no obvious relation to bond strength, instead of recognizing that it is the making and breaking of bonds that leads to enthalpy changes in reactions, students memorize an algorithm: products minus reactants. If Sanderson's suggestion were adopted, the standard state would be the gas phase atoms that compose the compound. This concept is consistent and applies not only to enthalpy but also to Gibbs energy and entropy calculations. A more detailed analysis of this presentation of thermodynamics can be found in references 6, 7, 9, and 10.
References
1 Dewey, J. Dictionary of Education. ed. R. B. Winn. New York: Philosophical Library, 1959.
2 Johnstone, A. H. "Chemistry Teaching -- Science or Alchemy?" Journal of Chemical Education 74 (1997): 262-268.
3 Spencer, J. N. "New Directions in Teaching," Journal of Chemical Education 76 (1999): 566-569.
4 Farrell, J. J., R. S. Moog, and J. N. Spencer. "A Guided-Inquiry General Chemistry Course," Journal of Chemical Education 76 (1999): 570-574.
5 Gillespie, R. J., J. N. Spencer, and R. S. Moog. "Demystifying Introductory Chemistry," Journal of Chemical Education 73 (1996): 617-622.
6 Spencer, J. N., G. M. Bodner, and L. H. Rickard. Chemistry: Structure and Dynamics. 2nd ed. New York: John Wiley & Sons, Inc., 2002.
7 Moog, R. S. and J. J. Farrell. Chemistry: A Guided Inquiry. 2nd ed. New York: John Wiley & Sons, Inc., 2002.
8 Sproul, G. "Electronegativity and Bond Type: Predicting Bond Type," Journal of Chemical Education 78 (2001): 387-390.
9 Sanderson, R. T. Polar Covalence. New York: Academic, 1983: 30-32.
10 Spencer, J. N., R. S. Moog, and R. J. Gillespie. "Demystifying Introductory Chemistry: 1. Electron Configurations from Experiment," Journal of Chemical Education 73 (1996): 631-636.
James N. Spencer is the William G. and Elizabeth R. Simeral Professor of Chemistry at Franklin & Marshall College. He is co-author of a general chemistry text and has written more than 100 articles on research and education. He was chair of the American Chemical Society Task Force on the General Chemistry Curriculum, and a member of the Wisconsin NSF New Traditions Project on teaching. He has been honored with the 1987 Chemical Manufacturer's Association National Teaching Award, the Lindback Teaching Award, the Dewey Award for Research, the 1999 American Chemical Society National Award for Research at an Undergraduate Institution, and the 2000 Mid-Atlantic ACS Emmet Reid Teaching Award. He is currently a member of the AP Chemistry Development Committee.
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