Chemical Education Today
Commentary
Let Us Give Lewis Acid–Base Theory the Priority It Deserves by Alan A. Shaffer
Why is the Lewis concept so often overlooked at the introductory level? Custom and historical perspective account for the traditional emphasis on the Arrhenius and Brønsted– Lowry concepts provided in textbook acid–base reaction chapters. These are certainly valuable concepts for inclusion in the teaching curriculum, but by themselves fall short of explaining how and why polar covalent reactions occur, which the Lewis concept can do. Any concern over confusing students by providing an alternate perspective seems inconsistent with the expectation in these same texts that students can deal with the far more challenging concepts of wave– particle duality and standing matter waves. In fact, applying the Lewis concept provides an opportunity for the student to use previous course concepts to predict and understand reaction behavior, increasing the students’ confidence and retention of key concepts. To highlight this concept sequence, one of the most important recurring themes in introductory chemistry is that of structure–property relationships of atoms, ions, and molecules. A large part of a given text’s coverage is devoted to leading the student through the methods of predicting the physical properties of substances. The general sequence takes the student through atomic structural theory, electronic configuration, periodicity, chemical bonding models (ionic, covalent, metallic), and more specifically for covalent compounds, the skills of drawing Lewis electron-dot structures and from this, the use of valence-shell electron-pair repulsion (VSEPR) theory to predict resultant properties related to polarity. Through this sequence of concepts, the central role of valence electrons is typically well emphasized and revisited, at first in the context of the chemical bonding models, and later with the finer details of molecular structure in VSEPR. With these skills in hand, students are ready to apply these to the realm of chemical properties and the “how” of chemical change, the reaction mechanism, at least for those reactions involving a simple change in covalent structure. In most texts, both at the full-year or one-semester level, the first major class of reaction to be studied in any detail is typically aqueous acid–base reactions, and the timing of this chapter sequence often closely follows that of molecular structure and properties. The fundamental reaction examples presented involve simple changes in covalent structure. These include the complete ionization of a strong acid in water, the net-ionic equation for neutralization of a strong acid and strong base,
the autoionization of water,
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and the reaction of ammonia with water to generate ammonium and hydroxide ions,
to mention a few of the most important examples. At this juncture, however, a major change in emphasis occurs, with something strikingly absent in the discussion of these reactions, as based on the survey of 22 introductorylevel textbooks conducted by this author.1 What could be called a “proton fixation” sets in at the very beginning of the chapter, with the critical role of valence electrons in causing these proton transfers being overlooked, while authors concentrate on the Arrhenius and Brønsted–Lowry concepts. At this critical juncture of first-study of reactions, the details of molecular structure so carefully crafted up to this point in time are largely removed, molecules are often represented only in script form (vs structural form), and valence electrons, so stressed before as essential in chemical bonding, are ignored. The student could get the impression that these protons just flip from one molecule to another, never recognizing that proton transfer is the structural result of electron pair displacement (1). The central and unifying concept of valence electron pair reactivity providing the how and why of proton transfer is at worst never addressed in many introductory texts, and usually, at best, discussed as an afterthought in a later section of the acid–base chapter in full-year texts. Authors stress the “what” of proton transfer, while ignoring or deferring the “how” of electron-pair displacement, and the “why” of nucleophilicity and electrophilicity. As an example, illustrations in these texts of the reaction of ammonia with water often show the proton transfer arrow, with or without the unshared electron pair on the nitrogen atom, depending on the author (see below). Some authors even allude to the necessity of this unshared electron pair in the reaction to accept the proton as a Brønsted–Lowry base, but the role of the electron pair in reaction is certainly not emphasized if not overlooked completely.
Of the 22 surveyed texts, only two alluded to Lewis acid– base theory with its focus on electron pairs, but neither developed the concept any further, though one had complete Lewis structures provided. Another text indicated that it would be “awkward” to apply Lewis’ theory to conventional acid–base reactions. This is an unfortunate perspective. Rather than an awkward hurdle to overcome, the Lewis concept empowers, and can be applied in a simple way to any conventional acid– base reaction encountered. An excerpt from a handout tutorial
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I have used in my general education chemistry course to provide supplemental coverage of this topic has been reproduced in the box on the following pages. Note how the arguments used in this tutorial serve both as a review and fresh application of previously covered course concepts to enhance the students’ skill at predicting chemical reactivity. This tutorial focuses on the ammonia molecule and its acid–base equilibrium in water. This analysis can be applied to any other acid–base reaction typically presented in an introductory text, some examples of which are shown below.
Of course, it should be stressed that the Lewis concept also applies to reactions not involving proton transfer as in the last reaction from the tutorial. The depth of coverage presented beyond the tutorial analysis can easily be left to the discretion of the instructor. Using this tutorial example can provide a lead-in to more involved discussion, including the concepts of nucleophile–electrophile or frontier orbital interactions, depending on the level of student. The complete skill of drawing electron-pushing mechanisms for polar covalent reactions is certainly appropriately deferred to the discipline of organic chemistry. However, this early exposure to the concept allows students to connect details of molecular polarity with reaction behavior. This connection is vital for a better understanding of the vast majority of polar reactions, including most organic reactions. It would be reasonable to share a few other examples with students, to give them the ability to predict the outcome of a given reaction. For example, in my classes I have been able, with practice, to get my students to push electrons correctly for simple reactions and predict the outcome of reactions like that of hydroxide anion with chloromethane, as illustrated in the tutorial. Students generally do value this early exposure to Lewis theory. As part of my introductory-level chemistry course (CH 101) end-of-semester survey from Spring 2004, I posed the question: How much did the inclusion of the Lewis theory of acids and bases help with your overall understanding of course
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concepts and principles? Did it tie things together or was it a hindrance? Do you think it is important to include in CH 101? Why or why not?
Of 33 responses, 27 were in favor of the value of the Lewis theory in helping with overall comprehension of course concepts and that course inclusion is therefore important. Six responses were of the opposite opinion. Of the favorable responses received, that best summarizing my students’ thoughts follows:2 The inclusion of the Lewis theory of acids and bases helped me to understand the course concepts and principles because it is a simple theory that summarizes most of all the chemical reactions seen in class. It helped me tie things together because all the reactions (or almost) are related to electron transfers [displacement]. And it’s important to include it in CH 101 because I think it is one of the basic principles students have to know in order to understand more complex chemical reactions. Once you know the Lewis theory it is easier then to fully understand other principles.
It is not so much a matter of rigor, or even what is most correct theoretically, but that Lewis acid–base theory, if taught well, ties more ideas together for our students, is most consistent with the proper emphasis of valence electrons on reactivity and so gives our students the best handle for the future study of new reactions. Indeed, electron-pushing mechanisms are the “algebra of chemical reactions” to quote one of my colleagues.3 An extremely valuable bonus would go to organic professors who would not have to unteach “proton-pushing” as a way of showing mechanisms, which most beginning and some veteran organic students use tenaciously, at least in my experience. We can not really blame our students for doing this— after all, “proton-pushing” is much of what they have seen from general chemistry and is, unfortunately, the major or only way their first exposure to reaction chemistry, in the form of acid–base reactions, is presented to them. There are exceptions to this oversight at the introductory level, for example the fourth edition of Ralph Burns’s Fundamentals of Chemistry (2). Burns brings in the Lewis theory after covering the Arrhenius and Brønsted–Lowry definitions, and he stresses the Lewis theory’s greatest versatility and universal scope. (Introductory texts could be easily revised to include a similar subsection within the acid–base chapter, perhaps revisiting the ammonia–water reaction from the Lewis perspective.) In whatever fashion, inclusion of Lewis theory at the introductory level needs to become the standard, not the exception. It is time again for us as educators to empower our students to see the more complete picture of chemical reactivity provided by Lewis’ theory.4 Notes 1. In October 2004 I used the Journal of Chemical Education Resource Shelf, Introductory Courses—http://www.umsl.edu/~chemist/cgi-test/mybooks.pl?category=26 (accessed Oct 2006)—to identify the
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Commentary 22 texts used in the survey. I thank Leslie Berger, Periodicals and Electronic Resources Librarian, East Stroudsburg University, for arranging the interlibrary loans that made this textbook survey possible. 2. Anonymous student course survey response, East Stroudsburg University.
3. Jones-Wilson, T. Michelle, East Stroudsburg University, private communication. 4. The author thanks Camille Law, Class of 2008, East Stroudsburg University, for her assistance in preparing the computer graphic files for this submission.
Student Tutorial Let us start with three major premises of reactivity based upon the concepts of stability and its relation to concentration of charge: • Negative charge seeks positive charge • Concentration of charge lessens stability and increases reactivity while dissipation of charge increases stability and lessens reactivity
Fine—but how and why does this reaction happen? We go back to the Lewis structure and use VSEPR theory to determine that ammonia has a triangular pyramidal molecular shape, has one unshared electron pair, and is therefore a polar molecule.
• Nature strives for increased stability
You are on a mountain ridge hiking with your friends, and you see dark clouds rolling in over another mountain ridge across the valley. You see a lightning strike and you begin to feel the hair on your head standing up. What do you do? You are about to get hit, and you do NOT have much time. Lay flat on the ground as fast as you can. This dissipates charge (it is spread out over your body and not concentrated on the top of your head) and makes you less of a target. See the tie-in with the premises above? Interestingly, these simple ideas can help us understand and predict much of nature’s reaction chemistry. Most introductory texts emphasize that acids donate protons and bases accept them. This rule is a good way to recognize bond making and bond breaking that occur with proton transfer reactions, and centers on the Brønsted– Lowry concept of acid and base behavior. Having said this, however, there is a deeper, more fundamental and more universal way to understand acid–base behavior. It has to do with the actual reaction chemistry going on (the mechanism of the reaction) and the changes in structure that result. What accounts for “normal” chemical reactivity (separate from nuclear chemistry) that we have focused on throughout the course? Is it protons? No! It is electrons and specifically the relatively high-energy valence electrons of an atom or molecule. If we look a little closer, we will see that electrons are responsible for proton transfer! To see this, let us focus on one specific example. You will recall that ammonia (NH3) is described as a weak base. It accepts to only a slight extent a proton from water to give the ammonium cation (NH4⫹) with water converted to hydroxide anion (OH⫺). The [OH⫺] increases and the pH rises above 7.0 since the [H3O⫹] decreases. The overall reaction is shown below:
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Now, what part of this molecule will be most likely to react with something else? Will it be the shared electron pairs already held between two nuclei, or the unshared electron pair sticking out by its lonesome in space? Hint: where is the greatest concentration of negative charge and therefore the site of greatest reactivity in the molecule? The unshared electron pair is this site!
Now we have identified the site of greatest negative charge concentration in the ammonia molecule. Negative charge goes after positive charge. Where is the site of greatest partial positive charge concentration in the reaction under study? It is equally either one of the hydrogen atoms of the water molecule because of the large difference in electronegativities of hydrogen versus oxygen.
Now we are ready to connect partial negative with partial positive charge to form a new bond. Bond breaking has to happen also because hydrogen cannot have more than two electrons. Electron-pushing arrows are used to show the role of changes happening with the valence electrons that account for new bonds formed and old bonds broken in the course of a reaction. Note how the unshared electron pair marks the tail of the arrow that depicts the “lone to bond” pair role change for the N lone pair; this arrow terminates at the H atom, which becomes the partner in the new N– H bond (Arrow 1 in the diagram below). Negative charge and positive charge are seeking each other. In concert with
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Literature Cited
2. Burns, Ralph A. Fundamentals of Chemistry, 4th ed., Prentice Hall: Upper Saddle River, NJ, 2003, pp 485–486.
1. Loudon, C. Marc. Electron-Pair Displacement Reactions, In Organic Chemistry, 3rd ed., Benjamin Cummings: Redwood City, CA, 1995, pp 90–91.
Alan A. Shaffer is in the Department of Chemistry, East Stroudsburg University, 200 Prospect Street, East Stroudsburg, PA 18301;
[email protected] Student Tutorial, continued this electron flow, the H–O bond pair evolves into an oxygen lone pair shown by the arrow from H to O (Arrow 2 in the diagram below). Carefully examine how these arrows account for the structural changes observed. Arrow 1 accounts for the new N–H bond in the ammonium cation, and arrow 2 accounts for the new lone pair in the hydroxide anion resulting from collapse of the pre-existing bond pair, which can be understood based on oxygen’s relatively high electronegativity.
Notice that in the process of ammonia accepting a proton as a Brønsted–Lowry base it has donated an electron pair to the proton of the water molecule. The water molecule in the process of donating the proton as a Brønsted–Lowry acid has accepted this electron pair from the base at the site of the proton, leading in turn to collapse of the original O–H bonding pair to oxygen. We do not always have proton transfers in reaction, but we very often have, excluding nuclear reactions, donation and acceptance of electron pairs as shown above. The descriptions above in terms of electron donation and acceptance are therefore more useful and universal. Gilbert N. Lewis, the same gentleman of Lewis electron-dot structure fame, formally proposed these definitions. A mnemonic (memory device) I use is connecting the underlined “a’s” below for acid. Then the base is opposite. An acid is an electron pair acceptor (thus Lewis acid in his honor). A base is an electron pair donor (thus Lewis base in his honor).
are used because what is forward or reverse is a matter of perspective. Note how electron-pushing arrows are used on both sides to account for changes occurring from left to right and from right to left. The role a molecule plays below relates to the reaction direction. (To save space below, “BL” stands for Brønsted–Lowry.)
Finally, note that the two directional arrows between left and right above are not of the same magnitude. This reflects that ammonia is a weak base, only undergoing this reaction with water to a slight extent. The reverse reaction (right to left) is much more favored. Though this has been a detailed examination with some new terms, the underlying principles of chemical reactivity are fairly simple and keep recurring. Look for this pattern in any new situation, and with a little practice you can predict what will attack what! This perspective will also give you a more complete understanding of many chemical reactions. For those of you who go on to take full-year general chemistry and perhaps organic chemistry after that, this perspective will be very helpful since most organic reactions can be described as a Lewis base attacking a Lewis acid. As an example, study the reaction below carefully, and you should be able to do some electron pushing to account for the observed chemical changes. Note also that no proton transfer is happening here but, knowing the products in advance, the same Lewis analysis can be applied. Indeed, this is a typical and important example from organic chemistry.
To complete the picture, we can make the following identifications for the reaction in the “forward” direction, that is from left to right, and we can also do the same for the “reverse” direction of reaction, from right back to left. Quotes
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