edited by: DAVID A. PHILLIPS Wabash College Crawfordsville,IN 47933 PRUDENCE PHILLIPS CrawfordsvilleHigh Schwl Crawfordsville.IN 47933
Why Teach Solution Equilibrium?
The Goals of Teaching Solution Equilibrium
James N. Butler Division of Applied Sciences Haward Univer~iN Cambridge. MA 02138
Paul T. Ruda Cleveland Hill U.F.S.D. at CheeMowaga Mapleview Drive Cheektawaga. NY 14255
Calculations of pH in solutions of acids and bases, calculations of the soluhilitv of minerals and orecioitates. and . . calrulations of the extent of complex formation between orranic lirands and metal ions are all Dart of the traditional curricuium of inorganic and physic2 chemistry. More recently, they have been adopted by those teaching geochemistry and environmental chemistry because of their relevance to the processes taking place in the weathering of rocks, the transformations of natural waters, and the fate of pollutants in the aquatic environment. Students read in the newspapers that "the pH of rain is normally 5.6" and in their textbooks that "the pH of pure water is 7.0." Why the difference? They read that acid rain of DH4.0 leaches toxic heavy metals out of soil and rocks and in& streams and lakes. t ow hoes this happen? They read that lime is heing added to Adirondack lakes to counteract the acid rain. On one level this is simple neutralization of acid by base. But how much to add? How long will the treatment he expected to last? What effect will it have on the mobilized trace metals? Would these mobilized trace metals he removed from drinking water supplies by the aluminum sulfate added for coagulation before filtration? At the high school and college general chemistry level, the curriculum tends to he limited to the simple proportions of stoichiometm and to the reometric diaerams and models of molecular structures, hutihe quantitative understanding of solution eauilibria reauires some facilitv with aleebra. Nothine more than quadratic equations is required in most cases, hut the leneth and comolexitv of mani~ulationsoften discouraee anyone who is not well-practiced in these matters. What then are the most important aspects of solution equilihrium to teach a t the elementary level? I would select the concept of pH as a master variahle controlling acid-base equilibria. Students are introduced to pH as a surrogate for "acidity" or "basicity," hut an even more important aspect is often neglected or poorly understood. This is the idea that if one knows the DH of a solution. then the ratio of acid to conjugate base €0; all acidbase pairs in that solution is also known. Such a eeneral theorem is easv to Drove: it follows from the ionization equilibrium for a weak acidbase pair:
My students are involved in a course of study that follows the New York State Regents Chemistry Curriculum, which groups kinetics and equilihrium into one unit. In 1981 the Reeents Chemistm svllabus was revised to allow the classroom te&her to reorga&'and emphasize topics of his or her choice. Therefore in my class, I generally treat solution equilihrium as a transition topic. This offers many excellent opportunities t o tie toeether the other tooics that I Dresent. One G m y goals in teaching solution equilibrium is to review the eeneral c o n c e ~of t eauilihrium and to make sure that the students rememher the characteristics of an equilihrium. We approach equilihrium from the macroscopic viewpoint as a steady-state system and a dynamic-reversible system a t the microscodc level. I t is i m ~ o r t a n for t students to he able to recogniz; a steady-state sjstem hased on their observations, whether aided lw instruments or not. The conceot of a closed system is critical to their understanding of theiteady state. Phase eauilihrium is introduced earlv in the curriculum. At that time I take a large, clear plastic jar, fill i t part way with water, cover it, and leave it on the window sill. The varied conditiona that occur allow the students to observe a variety of steady states. The iar is a source of many discussions and serves a centerpoint of the review. In my unit on water and solutions, I use the concepts of solution equilihrium in two ways. The first is the equilihrium of a gas dissolved in a liquid with the gas above the liquid in a closed system. Carbonated heverages are an excellent example of this type of system. Several articles are available that are written a t the level that the students can understand and enjoy ( I , 2). The students always enjoy discussing their experiences with carbonated heverages and the practical side of chemistry helps them understand why these events occurred. The second way in which I use solution equilihrium a t this unit is in a discussion of a closed system between a saturated solution and an excess of the solute. There are many examples of solution that can he used to show solubility and the various stages of saturation. I often carry out several crystal-growing demonstrations, which are not only very beautiful and fun to do. hut are also illustrations of the concepts heing prwnted. There is aLw a hands-on demonrtration that works auite well to aid student understandine . - .(3). . The effe& of temperature, pressure, and concentration chanees on eauilihrium svstems are also discussed. As this is donerthe importance of ;closed system is again stressed. The
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[HA1 - [H+l [A-I
K.
All too often, however, this equilihrium is presented in the context of the "dissociation of a weak acid," and the pH of a solution of pure HA is calculated as the first example. Historically, that may have been the first application of the concept, hut in contemporary life the master variahle concept is far more useful and important.
' Butler, J. N., "Carbon Dioxide Equilibria and their Applications."
Addison-Wesley Publishing Co.. Reading. MA 1982. 784
Journal of Chemical Education
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a mgl make involvs ms selsnian at msteial to be a, and htims to bedevoted tosach topic. For each column in this re, er, a him school and a college teacher have been invited todiscus why they tee1 a p&imlar topic is impartant and how it conhibutes to hstudents'
understanding 01 chemistry.
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(Continued from Col. 1,page 784) The next thing in importance is some understanding of the interaction of carbon dioxide, various other acids and bases, and calcium carbonate. This simole . svstem is essential for understanding geochemistry, environmental chemistry, and water quality engineering, and calculations can become very comp1icated.l But with the idea of pH as a master variable, it is easy to show how Con is transformed to HC03- and then to C0s2- as pH increases; how the presence of bicarbonate ion buffers the pH of natural waters in the ranee from 5 to 10: hut how strong kids (i.e., acid rain) exert their effects at pH below 4.5 independent of the presence of dissolved C09. It mi6;ht not be obvious at first, but the of finding the pH and total carbonate concentration in equilibrium with solid CaC03and atmospheric COz at 10-3.5atm partialpressure is easier (because of terms that can he neglected in the complete equations) than the problem of finding the pH and total carbonate concentration in pure water saturated with CaC03. By a selective choice of prbblems, and with adequate guides in the statement of the problem regarding approximations to be made, even a relatively inexperienced student can solve some of these very important calculations. Similarly, a selective choice of data and an explicit statement of approximations can give students a chance to calculate the concentration of lead eluted from a lead pipe by detergents containing complexing agents such as nitrilotriacetic acid (NTA). This too can be formulated as an eouilihrium r PhC03 is in equilibrium with'a solution problem. ~ nexample, of given pH containing small concentration of STA. Fixing the pH greatly simplifies matters, since it also fixes the ratio of the liaand I.'{- to the total dissolved liaand. It also fixes the concentration of PhZ+ in equilibrium witK the precipitate. The final ster, involves the association constant for the comnlex PhL-, b i t is no more complicated than a classical dissociakon problem. In all these computations, the graphical methods developed by L. G. Sillen and co-workers have been of great help to me, and I am pleased to see logarithmic concentration diagrams for displaying acidbase equilihria becoming part of the conventional curriculum. With the availability of small computers, students may return to algebraic methods more than they have in the past, hut for most the algebra will he clearer if they can see the whole picture as a graphic display. My main point has been that solution equilihria are an integral part of several important fields of chemistry, and that with a presentation that emphasizes their importance and uses the simplest types of ralculation as examples, even elementary studenis can gain some mastery of the subject.
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(Continued from Col. 2, page 784) concept of solubility product can also he introduced at this time. Solution equilibrium is also my bridge to the unit on acids and bases. Students have heard and used the terms acids and bases, but are most often not aware of the implicationsof these terms. Usine the concept of the dissociation of water, and Arrhenius' model for acidbase systems, students are introduced to acidbase chemistry. The Bronsted-Lowry model allows for an expansion of the utility of the concept. The relative streneth of acidmaw, the concentrations, their solutions (using p ~ \ n i t s ) , and the concept of buffer systems are all introduced (4). Another of my goals is to show practical examples of solution equilibrium on the everyday lives of my students. I feel that it is important to show practical examples of each concept that we teach to allow us to aid the students who have still not fully developed their abstract thinking skills. They need concrete examples to allow them to understand. "Real world" examples of these concepts also provide an answer to the frustrated students, who cry "why am I studying this, I am never going to use it."The impor&nce of solutions and solution eauilihrium in body chemistry, most especially blood chemis&y, is of special interest to me, and my students always enjoy this topic. The effect of medications on these systems is a good source for discussion. The article by John A. Lott offers a good introduction to blood chemistry (5).It offers an excellent review of this system and ties together many of the concepu that were previously presented. A second application that my students easily relate to is water pollution and water purification. Recommended levels of dissolved nutrients to maintain a healthy plant and animal population, biological oxygen demand, acceptable levels of inorganic waste in industrial effluents. and technioues of revitalizine lakes and streams are of great interest to my students due to our proximity to Lake Erie. There are many demonstrations and experiments that work well to illustrate this point. In summarv. althoueh solution eauilihrium is not addressed as one speci&c topicUin the curriculum, the concepts that comnose it are used at numerous times throughout the year, tying other concepts together and providing;mportant real life an~licationsof chemistrv. I feel that the time I devote to it is well spent. ~
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Literature Cited
.",.",\.",",. (5) Lou, John A., "Hydmgen lorn in B l d - A
51.6, (1978).
James Newton Butler is Gordon McKay Professor of Applied Chemistry at Haward University. He serves on ti-s facuity of Geological Sciences, the Committee on Oceanography, and as a trustee of the Bermuda Biological Station. His research has ranged fmm corhbustion to surface chemistry to electrochemical energy conversion to environmental Science and policy. Butler is best known for his 1964 text "Solubility and pH Calculations" and is atiihor or matiiha of mwethan 100 other publications. He received his BS degree from Rensselaer Polytechnic Institute and his PhD in Chemlcal Physics from Haward University.
Case of Dynamic Stsbility." Chsmlatn,
Paul T. Ruda is chemistry teacher and chairman of the mathematics-computer science department at Cleveland Hill High Schwl in Cheektowaga, New York. He has been active in many science teachlng organizations, and has been a conbibuta to the Testsd~monsmtiom feature in THIS JOURNAL. Ruda received the Distinguished Science Teacher Award from the Westem New York Section ofthe ACS in 1979 and the Distinguished Science Teacher Award from lhe Calspan Club of Sigma Xi in 1980. He received his Bachelor's and Master's degree from the State University of New York at Buffalo.
Volume 61 Number 9
September 1984
785
