A Mechanical Demonstration of Approach to Equilibrium
An Effective Demonstration of the Behavior of Indicators and Biochemicals in Buffers
Dean F. Martin University of South Florida Tampa, Florida 33620 Tested by: Kenneth Lothrop Marshfield High School Marshfield, Massachusetts At times it is useful to have a mechanical demonstration that combines audience participation with analogy to demonstrate a principle. he four-beaker experiment is one example of this type of demonstration. Two 4-1graduated heakers, each containing 2 1of water are placed side by side on the lecture bench, and two volunteers from the audience are requested. One student is given a 600-ml beaker, and the other, a 100-ml beaker. At a signal each student dins the beaker into a 4-1 beaker. eets as much water as possible and pours it into the other4-l beaker. At each transfer. the volume is announced and the students plot the data (ob;iously the sums of the volumes is constant, allowing for spillage). Obviously, one heaker quickly fills and one decreases, but a limit is approached after 8-12 transfers as is dramatically presented. The system demonstrates, mechanically, approaches to equilibrium hy two first-order reactions. The large and small heakers are analogous to specific rate constants k l and kz and the volumes remaining after a given trial are analogous to concentrations. Fimre 1is a CODV of data dotted bv a student. showinn the f ?;he scattk in volume in the eft(^) and kight ( ~heakers. the noints is due to the limitations of the eraduations of the heakers and the author's vision.
Charles L. Lerman' Juniata College Huntingdon, Pennsylvania 16652 Tested by: Richard R. Doyle Denison University Granuille, Ohio 43023 As a teacher of hoth freshman general chemistry and introductory biochemistry, I get two chances to help each student understand the nature of buffer solutions and the compounds on which they are designed to act. I attempt to formulate the following dichotomy. On the one hand, buffer components are compounds whose structures are used by chemists to control the acidity @H) of solutions by dissolving them at controlled, relatively high concentrations. On the other hand, BrQnsted acids and bases which are dissolved a t relatively concentrations in these same solutions are a t the mercv of the buffers, and the concentrations of the various possiblespecies are cokrolled by the acidity of the solution. The following demonstration is carried out after an initial theoretical discussion of the relative conrentrations of an acid and its conjugate base in buffers at various pH's. I have found it to he considerably effective in promoting students' retention of these concepts, and I use it hoth in freshman chemistry and as a refresher in the first few days of biochemistry. An arrav of solutions in test tubes is laid out in racks. one set of fiveat each of three pH's: 0 (1M HCI), 4 (1M acetate buffer). and 10 (1 M carbonate buffer). A DH meter can be used t'o"set the stage," i.e., establish firm$ in the students' minds the nature of the three environments to he used. Then the compounds p-nitrophenol and sodium p-nitrophenoxide are shown, hoth as structures on the blackhoard and as aqueous solutions, clearly establishing their colors. The students are told that the pK. of the acid is 7.1, and (in a small class) are asked to predict the color produced when some p nitrophenol solution is added to each buffer. Some student can usually do this quickly and correctly, and explain the predictions clearly. The tests are then carried out. The whole process is then repeated with the sodium salt to show that the behavior is independent of which form is initially introduced. Titrations of the solutions can be performed with concentrated HC1 or NaOH to show the ease of interconversion of the forms when the c a ~ a c i t vof the buffer is exceeded. Subsequent usebf other selected indicators demonstrates other principles, as well as providing more drill in situations where student participation is feasible. Methyl red (HA red, A- vellow: pKn 5.4) shows that hoth forms can he colored. ~ h i is ; important because, for many students at the freshman
Submitted by:
'cf. Hambly, G. F., J. CHEM.EDUC., 52,519(1975).
Submitted by:
Present address: Haverford College, Haverford, PA 19041. 634 / Journal of Chemical Education
level, color is tantamount to detectahility. The completeness of the color change emphasizes the sharpness of the concentration changes. Bromcresol green (HA- yellow, A2- hlue; pK. 4.7) introduces two new factors. First, it is a simple case of a non-neutral acid. (There is a sulfonate group of lower pK, which produces no color change upon dissociation.) Second, the pK, is near enough to 4 that the test solution in p H 4 buffer is green, and can he explained as observable amounts of both the yellow and the blue species. That is, when p H pK,, then (HA-) (A2-). Crude titration with strong acid or base shows how difficult i t is to re-achieve the green state (narrow p H range where both species are visible). m-Cresol purple (HzA red, HA- yellow, A2- purple; pKel
-
-
= 1.5,pK.2 = 8.2) is a useful contrast to hromcresol green, in that the color at intermediate p H is due to a single species, not to a mixture. Throughout the demonstration, the word "indicator" is purposely avoided. At the end, the students are asked what one might call the various compounds being added a t low concentrations. Generally, they know immediately, hut the experience brings a new understanding of what an indicator is-a compound whose-sole responsibility is to he "pushed around" by the solution in which it is placed. A final word to biochemistry students on the similar acid-base behavior of biological molecules in body fluids (e.g., why acetic acid is always referred to as "acetate" by biochemists) completes the point.
Volume 53,Number 10, October 1976 / 635
