MCTP Introductory Physical Science at College Park Personal Journal CHEM 121/122 Tom O'Haver 8/31/94 First class meeting, in our assigned classroom - a small generic classroom with only a blackboard and standard student writing desks. Eventually ten students showed up: six elementary education majors, a business major, an English major, a health administration major, and a studio art major. We started by introducing ourselves and telling a little about our special interests. I gave a brief overview of the course and handed out the syllabus. I explained that the course would employ a range of activities and that we would sometimes meet in a different location for laboratory work and for computer work. I then posed a question for the class and asked each student in turn for their opinion: What do you think is the greatest problem facing the people of the earth today? I summarized their many ideas on the blackboard - these included: overpopulation, pollution, violence, greed, poverty de-humanizing effects of technology, etc. I then asked them to consider whether science might be expected to make some contribution to the solution or reduction of any of these problems. I asked them to give some examples from their own experience of how we depend on science and technology and asked them to consider if there is ever any down side to the extensive use of technology in modern society. In all of this, the students seem very eager to contribute their ideas - not a shy bunch at all. As an assignment, I asked the students to read chapter 1 of the textbook. ---------------------------------------------------------------- 9/2/94 One student brought in a newspaper clipping about a chemical spill that had occurred on the Beltway the previous day and asked about the chemical significance of the materials that were named in that article. I asked the class to consider whether a basic knowledge of chemistry would be helpful is estimating the potential danger of such as spill. We spent a few minutes discussing the article. The main activity for the day was a scientific vocabulary class activity based on the relationship between scientific and common ('street') meanings of words. The idea was to get the students thinking about the meaning of words and to begin to have a notion that many scientific terms also have common meanings. I wanted them to realize that the scientific meanings are typically more limited and precisely defined than but are nonetheless related to the common meanings. I had in mind several words that students with little of know scientific background might be expected to recognize - such as: science, chemistry, physics, energy, atom, molecule, element, compound, force, pressure, heat, temperature, reaction, acid, neutralize, organic, inorganic, valence, charge, mass. I brought in two dictionaries for reference, circulating one of them for students to read the dictionary definitions to the class for comparison to our own ideas of meanings. This activity worked out well and generated a lot of student input. In fact, we only "got through" about half of the words on my list ---------------------------------------------------------------- 9/7/94 One student discovered that there was a periodic table of the elements in the front of the textbook and asked a question about it, related to the assigned reading. I took the opportunity to ask them to name materials on that table that were familiar to them and in what context. I asked the to locate and identify any examples of elements (as opposed to compounds) that they could see in the classroom or out the window. The formal activity for the day was to begin the "Air We Breathe" module. I distributed the student handouts with the observation that this was NOT a test or exam, nor was it homework, but rather it was to be an in-class group activity. I told them that talking with one's classmates and asking questions was encouraged, but that I wanted their own personal written answers to the questions and problems on the handout. I passed out compasses (for drawing a scale drawing of the earth and its atmosphere). Work on this module was continued the next class meeting. ---------------------------------------------------------------- 9/9/94 I began by asking students to explain their answers to some of the questions on the "Air We Breathe" handout. The most challenging part for most of them were the quantitative aspects, measurements and units, The scale drawing problem showed that some students are still not comfortable with using the concepts ratio and proportion to solve practical problems. Eventually, we designed a strategy for solving the volume-of-the-atmosphere problem, but delayed the computations until we visited the computer lab next period The class was told that the next class meeting would be held in the Macintosh Instructional Facility, Computer Science Center 3332. ---------------------------------------------------------------- 9/12/94 Today we met in the Macintosh Instructional Facility. Most students seemed generally familiar with the computers and the operation of the mouse - no one wanted to go through the mouse-practice program. While waiting for late-coming students to find the computer room, I passed out computer account applications forms to the students that did not yet have accounts. Then we began with an introduction to the ClarisWorks spreadsheet. I used the overhead video projection system to demonstrate some basic operations and the students tried things out on their own machines. Each student worked on their own computer, but there was a considerable about of peer help from the more experienced to the less experienced students. The students then began to work through a spreadsheet construction assignment that I handed out and I circulated around the room to help with problem spots. The greatest trouble spots were not related to computer of software operation but were related to fundamental logic and math concepts - measurements, units, ratio and proportion, etc. I passed out a floppy disk to each student and showed them how to initialize it and save their spreadsheets on it, I told them that they could work on their spreadsheets as homework, because ClarisWorks available on all public Macintosh computer rooms across campus and passed out maps showing the locations. We plan to continue working on this assignment for at least one more period. ---------------------------------------------------------------- 9/14/94 Today we met for the second time in the Macintosh Instructional Facility to continue our spreadsheet work. Most of the students opened the files they had saved at the end of the previous class and picked up where they left off, but two of the students discovered that they had not properly saved their spreadsheet file on their floppy disk and had to start over. Again, the main problems were related to logic and basic math. Most students struggled with unit conversions and ratio and proportion but eventually made good progress. Several students were able to complete the dilution-by-the-atmosphere problem that we had started in class two periods before. Most of the students seemed surprised by the result, and one student confided that she never thought she would ever be able to "program" a computer to solve a problem such a "hard" problem (in fact, SHE has solved the problem - the computer was just a tool). I must figure out some way to access the effect, if any, that the spreadsheet computer activity had on students conceptions of problem solving and of the use of that technology. Overall, I feel that this activity was too demanding for these students and I will probably need to "tone down" the quantitative emphasis and focus more on the conceptual aspects. ---------------------------------------------------------------- 9/16/94 A lot happened today. An article in the morning's Post about a car accident on Route 95 caused by the driver inhaling "nitrous oxide" touched off a discussion of the applications of this most interesting substance. I challenged the students to explain how it was possible for people to breathe nitrous oxide (as an anaesthetic, for example), while "nitrogen oxides" are listed by the EPA as a air pollutant with a permissible limit of only 0.05 ppm! Doesn't nitrous oxide qualify as a nitrogen oxide? We then discussed the spreadsheet work done in the previous class. I asked the students if they felt that the use of a spreadsheet was helpful, harmful, or made no difference in working out complex multi-step "statement problems" such as the atmospheric dilution problem. Students unanimously agreed that the use of a spreadsheet was helpful. They volunteered with the following three advantages: reduces arithmetic errors; allows the variable to be changed to see what effect they have; and the spatial layout of the numbers on the page, with cells and columns labeled, made the problem clearer. Good observations, I thought. Four of the students who had completed the volume-of-the-atmosphere and the atmospheric dilution problems volunteered their results, which I wrote on the board. They were in substantial agreement. We were all amazed at the result, which seemed counter-intuitive to everyone, including myself. At last our laboratory was cleaned and ready for us to use, so we spent the second hour in the lab, learning how to use the electronic digital scales, weighing our textbooks (which happen to weigh almost exactly one kilogram), and starting the "Counting by Weighing" module. Good response to working in small groups. Much surprise in the "guess the number of beads" part that their final measurements were often very far from their initial estimates. I announced a short quiz for next class, after which we will continue in the lab. ---------------------------------------------------------------- 9/19/94 We began with a 20 minute quiz, then moved to the lab We completed the Counting by Weighing unit and the bead jewelery construction activity. This went very well, with the students working well in pairs and obviously enjoying the "easy" activity. As usual, I discovered a few rough spots in my handout which can be improved upon. By the end of the period, everyone had completed the experimental portion of the activity but most had yet to complete all the discussion questions on the handout. Next time, we will meet back in the classroom to allow the groups to complete their discussions and calculations. ---------------------------------------------------------------- 9/21/94 Returned the graded quiz and discussed the questions with the class for a few minutes. Most students did well - the average was about 80%. The most common difficulty is with the one question that required a basic mathematical operation - units conversion. We spent the rest of the time reviewing and discussing our results from the counting by weighing and bead jewelery construction activity and making the connection to chemistry more explicit. ---------------------------------------------------------------- 9/23/94 The first portion of the class was devoted to small group pencil-and-paper work based on assigned reading from the textbook. I posed several questions dealing with the formation and depletion of the stratospheric ozone layer, the relationship between the ozone layer and solar ultraviolet radiation, and the effect of chlorofluorocarbons on the ozone layer. I got quite good answers - they must be reading. One group asked some very good questions about the the nature of the chemical bond that lead to a discussion of the electronic structure of atoms. Most students seem to have a classical "mini-solar-system" mental model of atomic structure, with discrete electrons "orbiting" the nucleus. One student wanted to know if the shared electrons in a chemical bond travelled in a figure-8 around both nuclei. Neat idea! I'm not sure how far into quantum ideas we should go in this class. In connection with another question about the difference between the terms "allotrope", "isotope", and "isomer", I mentioned the newly discovered allotrope of carbon, buckminsterfullerene (a.k.a. buckeyballs), which one of the students had heard of before. No one seemed to know who Buckminster Fuller was, but they expressed interest in the idea. Then we did a hands-on molecular modeling exercise, using small plastic model kits. Working in pairs, the students constructed models of various molecules, given their empirical formulas and attempting to follow the normal bonding behavior as implied by the number of bonding holes in each model atom. (The molecules, in the order of construction, were: hydrogen, chlorine, hydrogen chloride, water, methane, chloroform, ethane, propane, butane, ethanol, oxygen, carbon dioxide, freon, and octane). They investigated the geometrical properties of the molecules and compared the structures to the conventional "flat" drawings in the textbook. One student volunteered that, as a result of this activity, she realized that many molecules were 3-dimensional and that the they are flexible and can flop around more than she had imagined from the usual textbook pictures. Students were challenged to discover the concept of structural isomers, by trying to find alternative structures for butane and for ethanol that has the same empirical formula but a different "connectivity". The difference between "morphology" and the "topology" of molecular structures emerged quite naturally from this activity (but not using those scary words). One group spontaneously discovered the possibility of "ring" compounds. A question that arises is: does any molecule that one can build following the common bonding rules (e.g. "octet" rule) actually exist, and conversely, does every molecule that actually exists follow the common bonding rules? We will gradually discover that the answer is NO, but that it does not diminish the usefulness of the the bonding rules. (Do the exceptions really prove the rule?) At this stage most of the students don't have a very clear idea of the ORIGIN of the bonding rules, in terms of atomic structure, but those ideas will gradually emerge with further experience. Students expressed quite a bit of interest in buckminster-fullerene, but their student model kits are too small for such a large molecule. So at the end of the class we gathered all the model kits together and two of the groups tried to construct a model - without success. We needed to learn more about that structure. This molecular modeling activity was *very* successful. Everyone had lots of fun. I plan to bring the model to class regularly so we can continue to develop our experience and intuition with chemical structures. Perhaps by the end of the term we will be able to construct models of proteins and even DNA, by combining several of the kits. ---------------------------------------------------------------- 9/26/94 Today we began a new laboratory experiment, dealing with temperature and its association with molecular motion. The students observed a range of phenomena that suggest that there is a connection between motion and temperature (friction heating, evaporation, boiling, evaporative cooling, condensation). These observation will eventually contribute to the hypothesis that molecular motion is the physical meaning of temperature. The students were asked to propose explanations of the underlying mechanisms of these phenomena, based on their prior knowledge and the observations they made in the laboratory. I was very pleased to see that these activities generated a great deal of questioning on the part of the students. In fact, this was an activity that generated many more questions than it answered. At the end of the period, many question were still not answered or even fully explored. For example, many students were puzzled by their observations of evaporative cooling and could not suggest a reasonable hypothesis for a molecular mechanism. The problem is that they have no reason to suspect a *distribution* of molecular velocities at a given temperature. At this point it would probably be a good idea to introduce a molecular kinetics simulation, either computer-based or a macroscopic physical model (e.g. pinballs in a vibrating frame). I have located a simple program for MD-DOS, called "Teddy", that simulates this nicely for one gas or a mixture of two gases of different molecular weight. ---------------------------------------------------------------- 9/28/94 State Consultant's Panel meeting. Class was met by the TA. ---------------------------------------------------------------- 9/30/94 Exam 1. Asked for questions before distributing exam and spent 30 min or so responding to questions. I consider this to be an "easy" exam. ---------------------------------------------------------------- 10/3/94 Having given our first "hour" exam during the last period, I returned the graded exams to the students and went over the entire exam question by question. Results were generally good, but some students are still week in some basics, such as units conversion and numerical prefixes (e.g., recognizing that tetra- means 4). An article in the paper about U. of Maryland engineering students winning a national competition with their "alternative fuel" car prompted one student to ask a question about automotive fuels, in particular the effect of the addition of alcohol and of the difference between diesel fuel and regular gasoline. Our textbook did not have much information about that question so we agreed to put that on hold until we could check more reference material. Most of the time was spent comparing and interpreting our results from the last lab experiment, in which we observed temperature, friction heating, evaporation, boiling, evaporative cooling, condensation, and thermal expansion. We tried to make sense of the apparent connection between motion and temperature and to come up with a model of what is going on at the microscopic level that might explain all these observations. I made the point that the chemists' view of such phenomena tends to seek the microscopic mechanisms underlying the macroscopic observables. There were some good "what if" questions from students, one wanting to know if water would boil if it were contained in a rigid container. Evaporative cooling continued to be the most challenging to explain. One student proposed an explanation based on the idea that the "faster molecules escape" leaving the slower ones behind. This prompted another student to wonder if it would be possible to "make water boil at room temperature" by reducing the pressure. I said that this was a very interesting idea and that we would try it in the lab the next class. (We have the equipment to do this at least for a more volatile, lower boiling liquid, if not for water, say, ether or acetone. It's the same principle.) ---------------------------------------------------------------- 10/5/94 We began by meeting in our laboratory to resolve experimentally a question raised by a student in the last class - namely, is it possible to make water boil at room temperature by reducing the air pressure over it? It didn't work for water, using a simple water aspirator as a source of vacuum. However, we then performed a modification of the experiment using a liquid with a lower boiling point than water, namely acetone, which was familiar to the students from its use in nail polish remover. This worked well - the acetone boils vigorously sitting right there in a flask on the benchtop at room temperature (20 degrees C). I asked the students to feel the bottom of the flask - some were afraid to touch it because they assumed that a boiling liquid must certainly be hot. Those who did muster up the courage to touch the flask were surprised to find that it was not hot - rather, it was ICE COLD! The temperature got down to as low as 6 degrees C, from the original 20 degrees. Much surprise and confusion. We talked about this at some length and I encouraged the students to recall their experience with or memory of pressure cookers and cooking at high altitudes, in order to establish the connection between boiling and pressure. One student finally came of with a reasonably satisfactory explanation, but I'm not sure he was successful in convincing all the other students. The next time we are together, I must ask some of the other students to provide their own explanation. After that we moved up to the Chemistry Computer Room, where we have access to the department's collection of chemistry software. There we began an activity using Brooks-Cole's Beaker, an interactive structure-drawing-cum-expert-system tool, and MacMolecule, a public domain 3D molecular visualization and animation tool from the University of Arizona. (The student handouts for this and other activities in this course have been posted on the MCTP gopher area, Inform.umd.edu, in the path: University of Maryland System and State of Maryland/University of Maryland System-Wide Projects and Resources/Maryland Collaborative for Teacher Preparation (MCTP)/Modules). Substantial variation in the pace with which the students progressed - we will meet in the computer room next time to continue our work. It strikes me that should be making a connection between the structure of molecules and the mathematics of "graph theory", but I don't know enough of the terminology and foundations of graph theory to make a stab at it. I must look at that section in Jim's math course more carefully. ---------------------------------------------------------------- 10/7/94 Students finished the Beaker/MacMolecule activities begun last class. These activities have the students working with electron configurations, Lewis Dot Diagrams, the "Octet Rule", predicting likely structures, learning to draw chemical structures using Beaker's specialized drawing tools, discovering alternative ways to display a given structure, including the widely-used line-segment shorthand notation, investigating Beaker's rule-driven IUPAC naming capabilities, and trying to fool the program by drawing a hydrocarbon structure that is too complicated for it to name. In the MacMolecule activity, student view and manipulate full-color animated models of molecules in wireframe, ball and stick, and the more realistic space filling representations, in order to develop their 3-dimensional thinking. Molecules they work with include the "recent" discoveries buckminsterfullerene (a.k.a. "buckyballs", a newly discovered form of carbon) and compounds of the rare gases, which were once thought to be inert. (Later in the course, we will return to this program to visualize and manipulate larger biochemical molecules such as amino acids, polypeptides, proteins, DNA, and some vitamins). Student are continuing to make progress in their ability to describe and predict chemical bonding and molecular geometry, and I continue to make progress in my gradual conversion to the constructivist philosophy. I also am learning, but only AFTER the students have gone through my activities, how to make them better. My most common mistake is jumping past too many unspoken assumptions and intermediate conceptual steps. Some of this is corrected in real time during the class, but I still feel I can do a better job in designing the activities. I am gratified to see that a few students are willing to "play" with ideas we are working with and to go beyond the explicitly stated requirements of the handouts. ---------------------------------------------------------------- 10/12/94 Someone once said that the difference between art and science is that the goal of art is to convert the ordinary into the extraordinary, by constructing extraordinary works from simple everyday components, whereas the goal of science is to convert the extraordinary onto the ordinary, by reducing and recasting incomprehensible observations into models based on simple principles. For us, the "incomprehensible observations" are provided by the behavior of room-temperature acetone under reduced atmospheric pressure: we all observed that the acetone boils at room temperature and simultaneously becomes colder and colder. We spent much time today trying to expand our model of heat and temperature to explain this puzzling observation. Based on our previous laboratory observations, we had adopted the notion that increasing temperature is associated with faster random molecular motion. But it was more difficult to explain the observation that acetone could be made to boil at room temperature by reducing the air pressure in the container, as well as the anecdotal evidence from cookbooks that eggs and pasta take longer to cook in Denver because the boiling point of water is lower there, and that food cooks faster in a pressure cooker than in an open pot because water can be heated to temperatures greater than 100¡C under pressure. To explain these observations we had to extend our model of boiling to include the role of the atmospheric pressure. As one student put it, the atmospheric pressure "squeezes the molecules together", counteracting the tendency of random molecular motions to pull the molecules apart. When the molecular motions pulling the molecules apart exceeds the atmospheric pressure pushing the molecules together, boiling occurs. This boiling can be observed either by increasing the temperature or by reducing the pressure. A possible explanation for the evaporative cooling effect was suggested by observing the behavior of a collection of small hard balls contained in a vessel connected to a mechanical vibrator. The vibration of the container walls is a model for molecular motion, which is rapidly transmitted to the balls. The balls bounce about wildly and randomly. As the the container walls are made to vibrate more strongly, simulating a higher temperature, the balls were seen to move faster. It was clearly seen that the balls were not all moving at the same speed - at any given instant, some balls were moving fairly slowly, others were moving more quickly, and a few were moving very fast. But in spite of this variation between the balls, it did seem that the AVERAGE speed of balls increased as the vibrator was turned up. Moreover, a few of the fastest balls occasionally achieved enough speed to fly out of the container, and the number of balls that reached that speed increased with "temperature" - a model for evaporation. If the fastest balls were allowed to leave (evaporate) then it stands to reason that the average speed (temperature) of the remaining balls would be lower. One student expressed the concern that her explanations were not "scientific" in the sense of using the official technical words. I assured the students that I much preferred explanations in their own words that make personal sense to them. ---------------------------------------------------------------- 10/14/94 At the end of the last class I had given the students a few pages from Harold McGee's book "On Food and Cooking" to read; a section that talks about cooking methods, heat transfer, browning reactions, how microwave ovens heat food, etc. (A great book, by the way). This generated quite a few questions today. It's good to see so much natural curiosity - and I do encourage it - even when it means we don't "cover the material" as quickly as we might. Fortunately, almost everything can be connected back to important fundamental principles. (Back when I used to teach the mainstream science-major freshman chemistry course, I never got that many questions about science - mostly I got questions about course policy). We discussed a previous reading assignment (on chemical reactions and energy) and then divided up into pairs to work on a handout dealing with bond energies, heat of reaction, and the chemical basis for the caloric values of different fuels. Again I notice how well the students seem to work in small groups, talking between themselves to work out problems and develop explanations. Some of the groups bogged down when it got to the calculations, and I spent a good bit of time helping individual groups. We will continue this work the next time we meet in the classroom, but the next class we plan to do another laboratory experiment, making observations of several chemical reactions that involve the release or the absorption of energy in the form of heat or light. The one thing that continues to bother me is absenteeism. This week I have been averaging only about 80% occupancy - not very good. ---------------------------------------------------------------- 10/17/94 We were back in the laboratory today, and again I see how much the students really enjoy these hands-on experiments. They produced and observed three chemical reactions: a spontaneous exothermic reaction that releases heat, a spontaneous endothermic reaction that causes the surrounding solution to become cold (a possibly surprising observation); and a reaction that releases light energy (a "lightstick"). We looked for evidence of reaction, measured temperature changes, observed mass changes (or the lack thereof), discovered the effect of temperature on the lightstick reaction and tried to explain it. So far we have not talked very deeply about what is going on at the molecular level, but it will be especially interesting to see as they struggle to explain these observations in molecular terms. I posted this activity, and others I have done this semester, on the MCTP area of the U. of MD. gopher, InforM.umd.edu, in the path: University of Maryland System and State of Maryland/ University of Maryland System-Wide Projects and Resources/ Maryland Collaborative for Teacher Preparation (MCTP)/ Modules/. I would like to get some other peoples' activities posted as well. ---------------------------------------------------------------- 10/19/94 Today we finished up a pencil-and paper activity called "Counting Bonds and Calories: A Molecular View of Reaction Energy", in which the student learn to predict reaction energy changes ("Heats of Reaction") by tallying the bond energies of reactants and products, and compare the calculated estimates to the experimentally measured values available in the literature. In this way they were able to predict that ethanol will have a lower heat of combustion that propane and other hydrocarbon, that heat could never be obtained by burning nitrogen, and that the measured heat of combustion of wood is close to the value they calculate for glucose (expected because wood is composed of a kind of cellulose, which is composed of chains of simple sugar units similar to glucose). The concept of comparing an calculated theoretical estimate to an experimentally measured value seems new to many students. In particular, some of the students were puzzled at first when their calculations did not agree perfectly with the literature values. Their first response was "What did I do wrong?". I talked with one of the groups about assumptions in the model and the idea that the experimental values must be considered the "correct" numbers, in spite of inevitable experimental error. Of course, there is no general answer to the question: how much disagreement between theory and experiment is too much? Students tend to like nice neat hard-and-fast rules, but they must learn that the real world is messier. In the second hour, we talked about the last lab experiment (Reactions and Energy). I asked students to describe and interpret what they had observed. The most questions and thoughts concerned the lightstick experiment (these being seen often in the stores now that Halloween is close). One student came up with the idea that the light was the result of photons, but then she decided that the chemical reaction must be absorbing photons. She was simply confused about the light being a product or a reactant in the reaction. Some students seemed surprised that something as immaterial as heat or light can be thought of as a part of a chemical reaction. I myself was surprised by how well the students explained the effect of temperature on the light intensity. I got some quite good explanations of the connection between temperature, average molecular motion, collision frequency between reactants, and reaction rate - not in those exact words, but the concepts seem to be there. ---------------------------------------------------------------- 10/21/94 Sometimes I find that basic concepts and skills can be stumbling blocks for many students. Today, for example, we were comparing the energy required to decompose water into its elements 2H O (g) --> 2H (g) + O (g) 2 2 2 to the energy required to vaporize water H O (l) --> H O (g) 2 2 in order to shed some light on the nature of the bonds between the atoms within the water molecules, compared to the bonds between the water molecules themselves. One student was very confused because she could see no difference between the two. It was only after I had the student attempt to draw a PICTURE of a collection of molecules representing these processes that she understood. In science we very often use symbolic notation and special words to represent physical processes that we picture in our minds. It's very easy to forget that students may need lots of practice in converting from the mental pictures to the symbols and words and vice versa. In another activity, we were attempting to perform a calculation that called for the conversion of a measurement from one set of units to another, given the relationship between the units. A few of the students are still having trouble with this. Some of them seem never to have been exposed to the idea of applying mathematical operations to the UNITS of real quantities - the kind of thing that used to be called the "factor-label" method or "dimensional analysis". Did this fall out of favor? I am wondering if I should have spent even more time on this in the beginning, perhaps in the form of a more detailed formal lesson. One student did say that they had just been performing similar operations in math class, so perhaps the problem is a reluctance to generalize and transfer between disclipines. ---------------------------------------------------------------- 10/26/94 Today we began a new laboratory experiment, based on the "Acids, Bases, and pH" module. I asked the students where they had heard of "pH" before; they mentioned "pH balanced" deodorants and shampoo. One student recalled a connection to swimming pools. I asked them if they knew what it meant, and there was hesitation. We began with a brief introduction to the pH meter, a simple electronic instrument for measuring pH, and the use of "pH indicator paper", small strips of paper that change color in response to pH changes. They measured the pH of a wide range of familiar substances, including water (distilled and tap), vinegar, lemon juice, orange juice, tomato juice, milk, carbonated water, Coke, 7-Up, sugar water, salt water, milk of magnesia, shampoo, household ammonia, wood ashes, lye, hydrochloric acid, and sodium hydroxide. We noticed some trends: all the food substances were slightly or strongly acid (pH < 7) and never basic (pH > 7), for example. There were some surprises: salt water (made from table salt) was noticeably basic (which I wouldn't expect for pure NaCl) - it may be due to the NaI (it was iodized salt) or to the anti-caking agent. We have not begun to attach a quantitative meaning to pH yet - we are only getting a feeling for the range of likely pH values of common substances. One student commented that the pH of orange juice (2.8) was surprisingly close to that of a solution of HCl (2). As they left, I gave each students a package of 100 pH test strips to take home and play with, with the charge to test the PH of the food they eat and drink during the day to see if food and drink really is always acid rather than basic. ---------------------------------------------------------------- 10/28/94 At the end of the last class, I gave each students a package of 100 pH test strips to take home and play with, with the charge to test the PH of the food they eat and drink during the day, to see if they notice any trends. The results are in, and it does indeed seem that food and drink, and least the food and drink typically consumed by our students, is always either acid or neutral, but never basic. This may be because the response of our taste to acid substances, usually described as "sour", is more pleasing that our taste response to basic substances, which is usually described as "bitter". We continued to gather experience concerning pH. We discovered that if we start with a water solution of an acid (pH < 7) and add a base in small increments, we can increase the pH until it is equal to that of pure water (7), and that further addition of base will further increase the pH into the basic region (pH > 7). Since the solvent for these solution is water, and since we, as water-oriented creatures, consider water to be "neutral", the process is logically called "neutralization". The word has, of course, some different but related meanings in everyday language. In an attempt to understand the meaning of pH in a more quantitative manner, we tested the effect of dilution of a solution of hydrochloric acid with water, and we discovered that the pH increases (becomes closer to that of water) as more water is added. This was not particularly surprising. Specifically, we started with a hydrochloric acid solution that had a concentration of 0.01 moles per liter and found that it had a pH very close to 2. Then, by diluting it by a factor of 10 (one volume of original acid solution plus 9 volumes of water), the concentration of the acid would be expected to be 0.001 moles per liter, and the pH of the diluted acid was found to be very close to 3. Is this a trend and would it continue? When another 10-fold dilution was attempted, however, the different laboratory groups began to get inconsistent results, so at this stage we have not reached a definite conclusion, It does seem, however, that pH has an "interesting" relationship to acid concentration, and that the measurements and manipulation of very dilute solutions can be challenging to reproduce with consistent results. As they say - it needs more work. ---------------------------------------------------------------- 10/31/94 Continuing to do pH work in the lab today, we ran into considerable difficulty with poor quantitative agreement between the results obtained by the different groups. In the HCl dilution series, most of the groups obtained good agreement on the measurements of the pH of the original (.01 molar) solution and the 10-fold diluted solution, averaging 2.02 and 2.98, respectively. However, the 100-fold dilution were quite variable. I asked the students to try a 2-fold dilution, and the groups reported good agreement (averaging about 2.35). So it seems that it is mainly a matter of dilution - the least dilute (highest concentration) solutions give the most reproducible results. This is to be expected, as the more dilute solutions are more likely to be effected by contamination due to unclean glassware. In any case there was at least suggestive evidence that the pH increases by approximately one unit when the HCl (and therefore H+) concentration decreases by a factor of 10, as expected from the definition of pH. We obtained similar results with the sodium hydroxide dilution series; the two most concentrated solutions gave generally consistent results, but the more dilute solutions were wildly variable. Still, at least from the first two solutions, there was strong evidence that the pH decreases by one unit when the NaOH (and therefore OH-) concentration decreases by a factor of 10. Further trouble came when I attempted to work individually with two of the groups who were further along with their measurements, trying to get them to use algebraic manipulations to derive from the definition of pH an expression for H+ concentration as a function of pH (i.e. [H+] = 10^-pH). This caused some visible discomfort, as if recalling a painful memory. (One student said "You're hurting me now.") Perhaps it's understandable that students find it difficult to accept that pH was invented simply as a convenience, to avoiding having to write out very small hydrogen ion concentrations in scientific notation (e.g. saying 3 instead of 10^-3). To the students this seems a small advantage to trade for the "pain" of logs and exponents. ---------------------------------------------------------------- 11/2/94 Today we were back in the classroom to compare our laboratory results. The neutralization experiment (part 2 in the student handout) gave generally good results, with some groups showing a very nice s-shaped curve when they plotted pH vs drops of base added to an acid solution. (The theoretical treatment requires the solution of six simultaneous equations and leads to a cubic equation for the titration curve, but I must admit that I did not have the courage to work through the mathematics with this class, as I do with the chemistry majors). One student asked how it was possible for the pH to be finite when the quantity of added base was sufficient to completely neutralize the acid - he reasoned that if all of the acid's H+ ions reacted with the base's OH- ions, to produce water, there would be no H+ or OH- ions at all left in the solution, and the pH would be undefined. That, of course, is a very good question. The observations made by the students themselves, that the pH is indeed finite in neutral solution and in pure water, suggest that the reaction of H+ ions with OH- ions is not quite "complete" and that pure water ionizes by itself to a very small extent into H+ and OH- ions. By combining results from the different groups, we managed to summarise the results of the NaOH dilution series and attempted to discover a relationship between the OH- concentration (known by dilution of the original standard solution) and the H+ concentration (calculated from the measured pH). An attempt to plot OH- concentration vs H+ concentration, expected to yield a hyperbola, was complicated by the wide numerical range of the data and the small number of useful data points. However, when we calculated the PRODUCT of these two variables, we obtained a value of 10^-13.8 at 0.01 molar NaOH and 10^-13.9 at 0.001 molar NaOH, suggesting that this product is nearly constant over wide ranges of pH and in very good agreement with the "accepted" value of the ion product of water (10^-14 at room temperature). This is further support for the hypothesis of an "equilibrium" between water, H+ ions and OH- ions. It's clear that we "need more points" to make this more convincing; I'll modify the experiment accordingly. After this rather trying experience, the students were full of questions about various practical aspects of acids, bases, and pH, including the interpretation of claims for cosmetic and cleaning products seen on TV. They asked some questions I couldn't answer right away, but I promised I'd consult a reference and report back to them. Better yet, I'll bring a reference to class and we'll do it together. After class, one student stayed to talk for a few minutes about applications of chemistry to cosmetics; it was clear that that was much more interesting to her than the logical investigation of the fundamentals that we has just done. I think I will make up some extra reading material from Carl Snyder's book "The Extraordinary Chemistry of Everyday Things", which has an entire chapter on the chemistry of cosmetics. ---------------------------------------------------------------- 11/4/94 I brought in a copy of Carl Snyder's book "The Extraordinary Chemistry of Everyday Things" and read some sections to the class on the chemistry and cleaning action of soaps and detergents at the molecular level, surfactants and surface tension of water, the manufacture of soap (which would probably make a fun lab experiment), and the chemistry of sweat, deodorants, and anti-perspirants. This generated a remarkable degree of interest and spirited discussion. It seems that personal care products is a real hot button with these students. It seems more interesting to them that the environmental topics that have been the emphasis of the course. (Does this say something about our modern culture?) I gave the book to one student who was particularly interested in cosmetic chemistry, and she read the section on the protein chemistry of hair with great interest, stopping to read interesting bits out loud to the class dealing with the effect of pH on the optical reflectivity of hair cuticles (which gives hair its "shine") and the breaking of disulfide cross linking of keratin protein strands (which is how "permanent waves" work). It's nice to see all this interest in the practical applications of chemistry, but I worry that we should be spending this time on fundamentals and conceptual change. Next time we'll work on something more serious. ---------------------------------------------------------------- 11/7/94 This report was delayed by the pressures of the NSF PIs meeting and the National Panel of Visitors meeting last week. In a previous class, one of the student asked why bubbles and suds are formed by soaps and detergents. Simple enough, I thought. Our textbook did not have much to say about that, so I had promised to look into it in some other reference material. In my search, I ran across an interesting example of an apparent discrepancy in book explanations. Accordingly, I brought into this class two current books that attempted to explain soap bubbles. Both of the explanations are couched in terms of "surface tension", which our textbook did describe and which I amplified with a chalkboard drawing. (This would been a perfect time to let the students "float" paper clips and thumb tacks on the surface tension film of water, and then sink them by adding a drop of detergent, but I just didn't think of that in time - we'll try it later). Anyway, here are the two statements (verbatim) from the two references: James Birk, Chemistry, Houghton Mifflen, 1994. Page 399: "It is possible to blow quite large bubbles as long as the surface tension of the liquid is sufficiently high. The higher the surface tension, the stronger the film that is the bubble's surface, and the larger the bubble can be before it bursts". Dorling-Kindersly Science Encyclopedia, 1993. Page 128: "Soapy bubbles can be stretched into strange shapes because soap weakens the surface tension of water". I asked the students if they felt that these two explanations were contradictory. They all felt that they were. Then I asked them which one was the "correct" explanation, and exactly half of the students chose Birk's explanation and half chose the Dorling-Kindersly explanation! We know from common experience that agitating pure water will make bubbles, even without soap, but the bubbles are very short-lived. That soap weakens the surface tension of water can be demonstrated experimentally, for example by observing how small metal objects supported on a surface tension film are sunk by adding soap. So that suggest that pure water a too high a surface tension to form stable bubbles. On the other hand, bubbles need SOME surface tension to hold their shape - So what do you think? Are soap bubbles too complex a system to attempt to understand? The students readily accepted the notion that a sphere is the shape that has the smallest surface area for a given volume. Perhaps this is just something they have heard before. Then there is the idea that surface area would tend to be minimized by the intermolecular attractive forces - how can that be demonstrated or proved? Of course, a spherical drop of water has a much smaller surface area than a hollow sphere (bubble) that contains the same mass of water. So the bubble represents a sort of metastable state or local minimum, suggesting that there is an optimum range of surface tension for good bubble formation - too high and it collapses into a drop and too low and the bubble won't form in the first place or won't hold together if it does form. I obviously need to research this in greater detail. Anyone out there have any good references? ---------------------------------------------------------------- 11/14/94 Exam 2, taken in the previous class and returned to the students today, showed that there were some persistent "alternative conceptions" which have still not been displaced by instruction. Some of these suggest that additional laboratory experiments might help. For example, one student offered as evidence of the connection between temperature and microscopic molecular motion the observed temperature increase when the two reactants in a lightstick were mixed (part 3 of the "Reactions and Energy" experiment). He explained this by claiming that the temperature increase he observed resulted from the "excited molecular motion caused by shaking" the materials. Perhaps this notion could be tested by measuring the effect of manual shaking on the temperature of plain water. My prediction is that it would rule out the possibility of significant temperature increase by manual shaking. Another student, in explaining why water boils at a lower temperature at reduced atmospheric pressure, claimed that at the reduced pressure the molecules are more "relaxed" and that this causes more rapid molecular motion, facilitating boiling. Fearing the worst, I posed the following thought experiment to the class. Suppose you have two open containers on a hot plate, one containing a liquid A with a boiling point of 100 degrees and the other containing liquid B with a boiling point of 80 degrees. Liquid A Liquid B b.p. = 100 b.p. = 80 \ / \ / |-------| |-------| | | |o o | | | | o o | | | |o o o| \_______/ \_______/ +--------------------+ | Hot Plate | | set to 90 degrees | +--------------------+ What would happen if the hot plate were set to a temperature of about 90 degrees? Everyone agreed that the liquid B would boil and liquid A would not. Then I asked which liquid would have the fastest average molecular motion and thus the higher temperature, or would they be the same. The class was split on this issue. Several students claimed that liquid B must have higher average molecular motion because it is boiling and obviously moving more than liquid A which was not boiling. Most students seemed to have trouble with the notion that the temperature of liquid A should actually be expected to be higher than that of B, because B will stay at its boiling point as long as there is liquid present (neglecting the possibility of superheating). Most were willing to take the MACROscopic observation of gross liquid motion as evidence for MICROscopic motion, and thus predicted that the boiling liquid B would be hotter that the quiescent liquid A. Only a couple of students agreed that this experiment gives evidence for differences in intermolecular attractive forces - and that those forces must be larger for material A than for B. If these ideas are important enough, they would seem to call for more experimental observation (preferably with MBL equipment, right, John? ;-). ---------------------------------------------------------------- 11/16/94 The beginning of a new topic today: drugs and drug design. The students seemed relieved to be going on to something new. We began by discussing the previously-assigned reading about the development of aspirin, a designed chemical modification of a folk remedy (willow bark tea) known from the time of the ancient Greeks. I then passed out the molecular modeling kits and the students, working in small groups, built and described the geometrical properties of six-membered rings of -CH2- and -CH- units, discovering the flatness and rigidity of the "benzene" ring. They then construct models of acetic acid (the acid in vinegar) and salicylic acid (derived from the active ingredient in willow bark tea) and acted out a little animation in which these two molecules come together, breaking and forming bonds in such a way that aspirin and water are formed (the reaction used by the Baeyer Chemical company in Germany in the 1890's to produce the first commercial aspirin). I then passed around an old bottle of commercial aspirin tablets and asked students to smell them to see if there was any evidence of a reverse reaction - they all were very surprised to notice the distinct and recognizable smell of acetic acid (vinegar). By combining their kits and reforming into larger groups, the student constructed models of morphine and demerol, compared their structures for similarities, and looked up practical information about these drugs in the Physician's Desk Reference. We then turned to the idea of molecular symmetry and the possibility of right- and left-handed molecules. By constructing several small molecular models and their mirror images, and trying to superimpose the mirror pairs, the students developed a rule for predicting what structural features lead to molecular "handedness". We will develop this idea further in the next class. As they left, I gave each student a pair of plastic polarizing filters with the charge to take them home a play with them, singly and in combination, to investigate their optical properties and their interaction with various sources of direct and reflected light. I promised that this would be related in some way to molecular symmetry. There was very good interaction between the student today during the group work, with much evidence of "lights going on" over young heads, especially in the symmetry work. In an unrelated observation, during out mid-class break two students came to talk to me at length about science going on the other classes they were taking - nuclear fuel disposal in one case and the cosmological genesis of the elements in the other. Nice to see so much interest. The student in the astronomy class said that he was amazed at how their textbook could present such incredible and amazing cosmological theories in such a flat and matter-of-fact way, as if it were no big deal. Interesting comment. ---------------------------------------------------------------- 11/18/94 Optical isomerism (asymmetrical molecules with distinct left- and right-handed forms) provides a good opportunity to show some nice connections between mathematics (symmetry), chemistry (molecular structure), physics (polarized light), and molecular biology (enzyme-substrate interactions). At the end of the last class, I had given each student a pair of plastic polarizing filters and asked them to investigate their optical properties. In discussing with the class what they had discovered, one student stated there seemed to be a "grain" to the material (like the grain of wood, presumably), although there was no visible grain upon direct inspection. (This is because the "grain" exists at the molecular level). I directed the students to observe the daylight that was reflected from the floor and from the wooden desks. They seemed surprised to discover that reflected light is partially polarized. I tried to make some connections to popular culture (polaroid sun glasses and 3D movies), but it was clear that they had never really thought very much about how these things work. I asked them to explain what would happen if they tried to use polaroid sun glasses to watch a 3D movie, or to use 3D movie glasses as sun glasses, and how to tell if a pair of sun glasses in a store display is really polaroid or just neutral tinted. This produced quite a bit of discussion, some confusion, and eventually a few halting explanations. I ultimately used the "waving rope passing through a picket fence" analogy to provide a physical model. I passed some laser pointers (small low-power battery-operated diode lasers) around the class and asked them to determine whether the light from the laser was polarized or not (they still had the two polarizing filters). It was interesting to see the interaction between the students as they discovered how to do this (Do you use one or two filters? Rotate the filter or the laser? etc). I then asked the groups to inspect and investigate the polarization properties of some transparent objects, including pieces of glass and clear plastic and a series of clear liquids contained in glass bottles (water, sugar syrup, ethanol, isopropyl alcohol, glycerin, mineral oil). This was a rather chaotic and unstructured activity - perhaps too much so - in which various combinations of students experimented with various combinations of polarizing filters, lasers, and transparent objects. We did ultimately discover that there was some very interesting properties associated with the sugar syrup that was not shared by the other clear liquids, but it seemed unlikely that the students would discover on their own the idea that some substances could rotate the plane of polarized light, so I introduced this idea myself. (Did I give up too easily?) We ran out of time before we could connect this observation with the molecular symmetry of the materials in question. Next time..... In all of the activities, the students work from a handout that requires that they write down a description of their observations in their own words. However, for some students the verbal discussions and physical hand-on activities often runs ahead of the writing; I need to encourage them to sit down and write now and then. ---------------------------------------------------------------- 11/21/94 Continuing our work on molecular symmetry, I passed out the plastic molecular model kits and we constructed models for the various molecules whose polarization properties we had investigated in the last class. The objective was to gather evidence for the knowledge claim stated in the textbook that the ability of some substances to rotate the plane of polarizer to the left or to the right is associated with the existence of distinct left- and right-handed forms of the molecule. The groups were asked to determine which, if any, of the substances students had distinct left- and right-handed forms, if necessary constructing both the molecule and its mirror image. We had done such an exercise previously for simple four-atom molecules, but scaling this up to larger molecules proved challenging, especially for glucose. We adopted glucose as a simplified model for the structure of the corn syrup, the one substance the students had discovered exhibited the ability to rotate the plane of polarized light (the others were water, ethanol, isopropyl alcohol, and glycerin, a series of progressively more complex molecules all built, as is glucose, from a the same small molecular sub-units - mainly OH groups and methyl groups). The key question that came out in the student's discussion is what manipulations are "allowed" when trying to determine if two structures are identical or not. What does it mean to say that two molecules have the same structure? Clearly any bond-breaking is not allowable. Conversely, any whole-object rotation should certainly be allowable (although even this was accepted rather tentatively by some students). But after having seen the students struggle with this and discuss among themselves, I now appreciate how complex this problem is. In particular, molecules are often rather flexible ("floppy", as one student put it) and can be easily bent into an infinite number of "conformations" without breaking bonds. We had discovered that in previous experience with the model kits. There was quite a extended discussion among the student groups - and some disagreement - as to exactly how much manipulation one was "allowed" to do to try to get one structure visibly identical to another. Some students spent quite a while trying all sorts of twists and turns to get glucose and its mirror image to coincide before finally agreeing that they must be distinct structures. One thing I did notice is that one particular student, who had previously exhibited discomfort and some lack of skill in certain mathematical aspects, was quite skillful at manipulating three-dimensional models and visualizing their symmetry properties - and in convincing other students of his conclusions. As illustrations of the practical significance of optical isomerism, I handed out an article about thalidomide and then passed through the class some bottles containing small samples of the left- and right-handed forms of carvone, a fragrant terpene-like molecule. I asked them to see if they could notice any difference between the smell of these two forms. (One form smells like spearmint - its mirror image smells like caraway or rye bread). I discussed briefly which properties of left- and right-handed forms would be expected to be identical (e.g. bulk properties such as density and boiling temperature) and which might be different (those that depend on specific molecular interactions with other asymmetric receptor sites). Evidently there must be some left- and right-hand specific receptors in our noses. By this time a considerable difference in the rate of progress between the groups has become evident. I asked the faster groups to predict whether the degree of polarized light rotation in a transparent substance would be related to the path length through which the light passes and to design an experiment to test this prediction. One student, who had not been among the most attentive previously, was so intrigued by the symmetry properties of the molecular models that she began to build some other structures on her own. I challenged her to try to find the simplest molecule containing only carbon, hydrogen, and oxygen that has distinct left- and right-handed forms. She came up with one structure that I had not thought of and seemed quite delighted to have done that. At the end of the period, she asked if she might be allowed to borrow the molecular model kit over the Thanksgiving break. No harm in that! ------------- 11/28/94 Finishing up our work on molecular symmetry, I posed to the students the following challenge: what is the simplest molecule containing only carbon, hydrogen, and oxygen that has distinct right and left-handed forms? (By simplest, I mean the smallest number of atoms). I didn't really have an definite answer to this question, but I thought it would be fun and instructive to try. As each student group came up with a candidate, I asked them to defend their structure, that is, to explain to the class how their structure meets the stated requirements. If it met the requirements, I wrote that structure on the blackboard. Several proposed solutions did not pass that test. Ultimately, two groups came up with 11-atom solutions and two groups came up with 10-atom solutions. (When I tried this myself, I found only an 11-atom solution, so the kids did better than I did :-). There was much evidence of increased skill at visualizing and manipulating structures and their symmetry properties. We continued our explorations of the molecular structure of common drugs and hormones by constructing models of several steroids and describing the common structural features of all steroids. Each group chose their "favorite" steroid and constructed a model of it. I asked several students to explain to the class the significance of structural similarities and differences between some synthetic steroids and their natural analogs: namely, RU-486 vs progesterone and anabolic steroids vs testosterone (based on assigned reading). A piece in the news about the suspension of Washington Redskins fullback Frank Wycheck for alleged use of anabolic steroids provided a timely real-world connection. ------------- 11/30/94 We began a brief unit on nutrition with a simple series of activities that could easily have been done in a middle-school classroom, using as a source of data the nutritional labeling of the boxes of five popular breakfast cereals. We poured out a "typical" serving in a bowl and weighed it to compare to the serving size on the label (and discovered that the manufacturer's serving size is rather scanty, in the opinion of most students); computed the total weight of one serving that was not accounted for in the nutritional contents (carbohydrate, protein, fat, salt, vitamins and minerals) and discussed what the missing stuff might be; computed the percent fat and percent "fat-free", comparing it to other fat-free foods the students know of; computed the Recommended Daily Requirement for protein from the grams of protein and percent of the U. S. Recommended Daily Requirement (RDA) given on the package label. I asked each student to choose what they thought was the "healthiest" cereal and to defend their choice. Fun but not very challenging. These activities required only basic math skills, in particular percent, units conversion (milligrams/grams/ounces), and proportion. As I had seen before, these aspects still gave some students pause, especially computing the Recommended Daily Requirement for protein from the grams of protein and percent of the U. S. Recommended Daily Requirement. I asked one of the successful groups to explain this to the group having trouble with this. Some difficulties in classifying nutrient materials were evident - one student was not sure whether sugar content of a sweetened cereal should be included in the carbohydrate content and another wondered where the fiber content should be accounted for. In the next class we will investigate the structures of these materials and try to relate their structures to their properties. ------------- 12/02/94 Continuing our investigation into the chemical aspects of nutrition, each student constructed molecular models of "alpha" and "beta" glucose. Then, combining the models, we acted out the polymerization of those sub-units into chains, eliminating water. In this way we modeled the formation of starch or cellulose. I had hoped to be able to have the students discover the difference between starch and cellulose by constructing both "alpha" and "beta" linkages between the glucose units, but the structural difference proved to be too subtle to be very dramatic, at least with the model kits we were using. We performed a similar chain-building activity for the formation of "peptide" linkages between the students' amino acids sub-units, forming a polypeptide. We didn't attempt to build a complete protein. With a couple of the groups I discussed an interesting section in the book about the "sweetness triangle" and the relationship of molecular geometry and sweetness of Aspartame (NutraSweet). a dipeptide that is chemically not at all like any other sugar but is 200 times sweeter that sucrose. It was new to me - very neat. The students seemed quite familiar with the terms "saturated fat" and "unsaturated fat", and most had enough experience that they knew that saturated fats generally have higher melting points, but had not thought about the chemical meaning of "saturated" (saturated with WHAT?). We did not try to construct a complete fat (triglyceride) molecule, but we did build both saturated and unsaturated a fatty acids and tried to understand how the saturation of the molecule might effect the melting point. ------------- 12/05/94 Rather than begin a new unit on a new topic today, I decided to spend the last week of class in a some reflective activities that would seek to try to tie together and connect the array of topics that we worked with in this course. As the activity this class and next, I divided the class into groups of 3-4 students and set each group to the following tasks: 1. What do you think are the major themes or "big ideas" of chemistry, as embodied in this course? 2. With your partners, develop ten (10) questions, suitable for an examination, that test mastery of the main themes of this course. Of course, you must also provide the "right" answers and a scoring rubric that indicates how partial credit is to be assigned. Your team's grade on this activity will be based on how well your questions test understanding of the broad array of concepts that we dealt with in this course. Those questions that I judge "best" will appear on the final exam, so it is to your benefit to develop good questions. Questions that are "too hard" or "too easy", and questions that test simple factual recall will not be considered suitable. 3. "Performance accessment" is a term used to describe a type of test or examination question in which students perform some practical "real" task and is scored individually on the success of their work. The "real" task might mean a laboratory experiment or an activity involving manipulatives. Design a performance accessment that might be suitable for this class and which could be used to obtain an individual score. Take into account the practicalities of class location, equipment limitations, and time limitations (two hour total time limit for entire exam). The students warmed to this task slowly, but after a while they got into the swing of things and worked busily together, arguing among themselves about the relative merits of various topics and questions. In one exchange, a student suggested that an important concept was the distinction between element, compound, and mixture and proposed a question that asked students to define those terms. But another student in that group quickly injected that it would be better to design a question requiring students to USE an understanding of those concepts to solve a problem of some sort. Excellent! (Both of those students were elementary ed majors, by the way). We will continue this next class, when I will ask the groups to give example of their work. ------------- 12/07/94 I asked representatives from each group to choose two of their ten questions to write on the board for all to see and critique. We spent the entire period talking about these questions and asking students to propose answers. All of these questions, and the others that the students turned on paper in at the end of the period, tended to be very similar to questions that had been asked in previous experiments, quizes, exams and other work that students has done. They did tend to cover the whole array of topics that I had chosen to deal with in the course, but none of the questions were what I would call really original. A couple of good questions derived from informal verbal discussions that had not previously yielded questions (e.g., numbers 3e and 6 on the final exam). Overall, I found the question-generating activity to be a good alternative to the usual last-class review period, in that the students were much more actively involved. Also, we had some good discussions concerning what constitutes a good question, the difference between conceptual knowledge and direct recall, and what were the really important overall concepts in ths course.