Reaction energy profiles are useful to pull together topics that are widely separated in the typical chemistry textbook, and they play a major role in the discussion of reaction mechanisms in organic chemistry. Energy profiles may illustrate concepts that regularly give students trouble, e.g., enthalpy, potential energy, transition states and reaction rates. Described here is a quick demonstration relevant to reaction energy profiles for lecture or even a hands-on activity for students. Over the years it has been useful to perform early in the first chemistry course and then return to repeatedly when looking at additional details of a chemical reaction. It is particularly valuable for students in their first course or for those students who may need concrete examples of abstract processes. Preparation for the demonstration requires no solutions, and clean-up after class simply means removing the objects from the room.
Classroom discussions of a chemical reaction initially focus on the before and after (Scheme I) along with the abstract claim that molecules somehow store energy only to transfer it (∆H) during the course of a reaction. This activity helps students see the distinction between the stored potential energy, ∆H, and Ea, often a conceptual roadblock between thermodynamics and kinetics, and it provides students with a concrete feel for what goes on during the course of a chemical reaction. Too many students still see kinetics (Ea) and thermodynamics (specifically ∆H) as two perspectives on the same behavior. How many times have we encountered the preconception that exothermic reactions must be fast and endothermic reactions slow? And if molecules store a lot of potential energy shouldn’t the molecule react spontaneously and quickly? Ask students why molecules don’t react if they are storing so much energy. Answer: it’s all about Ea.
An energy profile for the reaction in Scheme 1 is shown above in Fig. 1. Among the reasons for the value of this activity is that students see the rearrangements within the “molecules” that accompany the reaction and directly see the results of the release of energy. The height of the pop of the hemisphere (see below) can be compared to the energy released when molecules of the reactants rearrange to form the product.
The demonstration uses four different hemispheres constructed from once useful racquetballs (Fig. 2) to distinguish between potential energy (stored energy within a molecule), reaction enthalpy, activation energy and one of the most daunting of many conceptual roadblocks, the transition state and activated complex — that which cannot be seen but must exist. The spheres are considered equal except for the location of the cut; the mass of the sphere is ignored. The actual sizes of the four hemispheres are not critical; it is only necessary that one is so short (D, 20 mm) (low Ea) that it typically will not remain inverted; a second one, slightly taller (C, 22 mm), remains inverted long enough to place it on the bench top and move away before it spontaneously “reacts”; and two others (A and B) are sufficiently tall (32 and 30 mm tall, respectively) that once inverted they remain so. The taller “molecules” store a great deal of potential energy, which students may directly experience when they physically invert the hemisphere in an endothermic “reaction” to start the activity.
Students are asked to think of the inversion of the hemisphere as a chemical reaction in which energy is stored in the molecule. The four “molecules” are storing energy (potential energy), and depending on Ea this energy may be readily released. Only the smallest hemisphere will likely pop on its own in an exothermic fashion, dramatically evident by the height achieved by the molecule. Students can even try to guess the height of the pops of each sphere.
Students can be challenged to think about the reactions of the other three hemispheres with more potential energy. Why don’t they react if they are storing so much energy? The spheres may be passed around and students encouraged to store the energy by inverting the hemisphere; students then should run the “reaction” by popping the hemisphere back into its normal shape for all four molecules. The amount of energy/effort they have to put in to pop the hemisphere back represents Ea.
Students should also be asked about the appearance of the “molecule” at the top of the energy curve. (See Fig. 1.) What shape is the hemispheric “molecule” likely to go through, and were they able to observe the shape? To go from reactant to product they may reasonably propose that the transition would be a flat disk, a reasonable shape but one not directly observed. They have just described the nature of a transition state, the highest energy state of a non-observable but reasonable arrangement of atoms (activated complex) as we go from reactant to product. Achieving this high energy, planar arrangement is what determines Ea and the rate of energy release for the exothermic reactions.
I often place the four inverted hemispheres on the bench and turn to the board to start sketching the reaction energy profile during which time one of the shorter hemispheres will spontaneously (∆G < 0, low Ea) pop ~3 m into the air. The height of the pop represents the energy originally stored in the molecule. For “C” (the slightly slower spontaneous reaction) it is also possible to trigger the reaction by dropping it flat (curve up) from various heights — this would be the activation energy. The higher you need to drop the inverted sphere to get it to pop, the more activation energy. The activation energies for both A and B are so high that dropping them from the third floor to our marble tiled atrium below did not trigger the reaction.
At the end of the activity it is useful to match actual examples of chemical reactions (e.g., oxidation of silanes, boranes, and hydrocarbons) with the hemispherical “molecules.” For example, “A” may represent the oxidation of octane in gasoline requiring a large spark and pressure to release the considerable stored energy; “B” may represent the oxidation of CH4 from natural gas triggered by a small spark, even a static discharge; “C” may represent the spontaneous, highly exothermic reaction between O2 and B2H6, once considered as a high energy fuel for rockets; and “D” may represent the spontaneous, highly exothermic reaction of SiH4 with O2.
Students can be asked to draw a reaction energy profile for each reaction above and match it to the behavior of each racquetball “molecule.” (Consider giving out these four reaction energy profiles and ask students to match.) They have four “molecules” with varying activation energies, potential energies, and expected enthalpy values. It is clear that ∆Hrxn, is not a predictor for the rate of the reactions; molecules with high potential energies need not react quickly.
There is a temperature component to the “racquetball reaction” but it is not sufficiently consistent to allow any numerical analysis. However, the two shorter “molecules” (C and D) do react more slowly when cold. To ensure a spontaneous reaction it is a good idea to flex C and D back and forth a few times before the demonstration. There are no particular safety issues for the activity, but students, or instructors displaying bravado, should not look down at the shorter hemispheres while waiting for them to pop up. Watching the instructor getting hit between the eyes may make the event more memorable for the students, but it does nothing to impress upon students the ever-present need for safety.