provoccrtive opinion Why Teach the Electron Configuration of the Elements as We Do? Roger C. Millikan University of California, Santa Barbara, CA 93106
"What's the use of it anyway?" That question once came from my daughter when she was struggling to learn the mysteries of lone division. I believe we should ask that same question often when nmsidering what to inrlude in a general chemistrv wurse. Mans t w i r s ha\,e become "standard," tu the point that every textbook seems to include them in just the same way. Yet even those topics should he questioned and reconsidered. Perhaps the diversity of goals of students in general chemistry could he served better by covering standard topics differently. The electronic configuration of the elements is such a topic. Most general chemistry textbooks devote one or two pages to an enormous tahle of numbers eivine the distribution of " electrons among the atomic orbitals. Often, for completeness, all the 106 known elements are listed. That tahle is then followed by pages of discussion of the regularities, irregularities, trends, suhshell filling sequence diagrams and memory aids aimed a t making that mass of numbers mean something. Generallv a set of rules is set forth (for examvle. half-filled d shells are especially stable) to explain thk irregularities. What's the use of it anyway? Those electron configurations are the wrong ones for most students to learn. They are for neutral, isolated, ground state atoms. How many students (or even chemists) ever work with isolated atoms? Sure, a few gas phase spectroscopists do, hut nearlv all eeoeral chemistrv ex~erimentsare done in water soluti"on. Nkarly all industrial chemistry is done in condensed ~ h a s e sNearlv . all oreanic chemistrv is done in solution. Why do we spend time and effort teaching beginning students that neutral chromium atoms have [Ar]3d54s' for an electron configuration? Perhaps we want to say that chromium has that configuration instead of [Ar]3d44s2because of the extra stability of half-filled d shells. But then later in the course we mention that in water solution Cr2+ is a stable ion with a 3d4 configuration whereas Cr+ with a 3d5 configuration is unstable. Naturally, students wonder why. What I wonder is, why we teach the configuration of isolated atoms in the first olace. The ~ufbarcor huilding up principle, supplemented with a few special rules for the order of suhshell filling, seems to give one the guidance needed for constructing a table of electron configurations. This is what many texts tell the reader. You can figure it out-just learn these rules. For many cases, it works. Unfortunatelv. " . for manv it does not. The Drocedure recommended by one text for figuring out electron configurations sounded to me almost like a comvuter program. So I tried to use the huilding up rules to write a-progrk that would t w e out the orbital confimration given the atomic number ;the element. When exceptions cropped up, I treated them as special cases in this manner:
10 IF N=29 T H E N P R I N T "Cr [Ar] 3d(5) 4s(l)" Out of 106 elements in the tahle of electron configurations, there were 29 special cases! Somehow, rules that work only 73% of the time seem hardly worth teaching. Indeed, my comvuter vroeram would have been shorter without enhod;ing the riles. One hundred and six statements would have done the job (IF N = l PRINT Is(1) . . .). In my experience, there is another regrettable result that comes from presenting students with a tahle of electron configurations early in their study of chemistry. After spending much effort learning the rules and configurations, students come to helieve in them. The fact that they apply only to isolated atoms is missed. The fact that electron configurations vary with the quantum environment comes as a shock later in the course. Even when the instructor is careful to point out these things, students tend to fall into the trap. So far, I have been carping and negative. Let me try to be more positive. I t is agreed that tables of electronic configurations of the elements do have their uses. Used in conjunction with periodic tables, they help solve many prohlems. So put them in an appendix for ready reference. In the same appendix. one mieht " also out a tahle of electronicconfieurations of stable ions. I would wager that ion configurations would show even more regularly in orbital electron distribution than that displayed by element configurations. I would continue to use element configurations to illustrate and justify the building up principle of orbital filling for many-electron atoms. But for a first dose, I would only go up to neon or perhaps argon. The complications of heavier atoms should be mentioned, hut there is no need to dwell upon them, nor to present a table of all 106 elements. I would also stress that interactions with the environment change orbital energies, spatial distrihutions, and occupation numbers. As part of the subject of electron configurations, I helieve one should display stable ion configurations. The dominant role of noble gas structures will therehv he em~hasized.This can be related hack to a ~ p r o x i mate sb~utioniof the wave equation if desired. In el&ronic configurations, nature is telling us what those solutions are like, after all. Lastly, I would refrain from putting forward general rules that do not work very well. Parts of chemistry are idiosvncratic. When that is so. we should admit it. Am I recommending we throw away the bottom two thirds of the periodic table? Not at all. The complications encountered in the heavier elements can and should he treated later in the course. Bv then the students will he readv for dealine with special cases. By then they will expect generk to have interestine excentions. Mv olea is for authors to resist the temptation to>ove;large am&& of data by throwing in a massive tahle that no one can digest. What's the use of it anyway?
statement
Volume 59
Number 9
September 1982
757
