Do We Need Orbitals?

Do we need to discuss orbitals in General Chemistry? They are a standard part of G-Chem textbooks. With the “Atoms First” approaches, orbitals show up quite early in the semester. In my G-Chem class, they typically show up at the beginning of Week 4, so my class has been immersed in them all of this week.

As a quantum mechanic, I enjoy talking about orbitals and other aspects of quantum chemistry that show up in G-Chem. This is a topic in which I am probably more knowledgeable than your average chemistry instructor. I can comfortably field student questions on the weird and wonderful world of electron behavior. However, I also see the pitfalls in the simplified presentation at the G-Chem level, often leading to student misconceptions that I then have to correct in physical chemistry and/or inorganic chemistry. After many years of teaching, I can state that the topics surrounding the notion of an orbital are for the most part confusing to G-Chem students. Do we need them at that level?

Why do we ‘think’ we need orbitals? Here are some things an instructor might say.

(1) Writing electron configurations is one of the key ‘outcomes’ in a G-Chem course. By that, I’m referring to the 1s² 2s² 2p⁶ 3s² variety, rather than the 2.8.2 Bohr model variety.

(2) Students need to know s and p orbitals so we can discuss hybridization, and this sets up a discussion of resonance/delocalized-pi systems and all sorts of language used in Organic Chemistry (the course that comes after G-Chem).

(3) How else will we cover that Molecular Orbital diagram in the G-Chem text book so we can show the superiority of MO theory and its prediction that O₂ is paramagnetic?

(4) The “simple” picture of electron sharing in covalent bonds is wrong. Students need to be able to describe bonds using orbital overlap.

(5) How else will you explain why the Periodic Table has the unique shape that it does (with the s-block, d-block, p-block, etc.)?

Here’s how much time I typically take on orbitals (and leading up to them), as represented by titles of each hour-long class period.

·      Waves and Photons

·      Interaction of Light & Matter, and the Bohr Model

·      Photoelectric Effect and Wave-Particle Duality

·      Quantum Numbers and Orbital Shapes/Sizes

·      Photoelectron Spectroscopy and Orbital Energies

·      Electron Configurations

Three other class periods that relate to orbitals are:

·      Periodic Trends

·      The Covalent Bond (introduction)

·      Hybridization

Most years I don’t cover the standard Molecular Orbital “mixing” Energy Diagram involving the 1s, 2s, 2p orbitals, but I do discuss the overlap of two 1s orbitals to form a bonding and anti-bonding molecular orbital, and relate this to the Pauli repulsion, when introducing the covalent bond. Most of the discussion in Periodic Trends uses the simpler shell model, and orbitals are explicitly invoked only to discuss the two kinks in the trend of increasing ionization energy across a row in the periodic table.

If I didn’t concern myself with orbitals, I would be able to dispense with a swath of material in my current G-Chem class and use that time to tackle other topics in-depth and perhaps introduce other relevant important concepts in chemistry. But first I must consider if they can be dispensed with. Let’s consider responses to the five listed objections.

(1) Do we need really need orbital-based electron configurations? Bohr’s shell model is sufficient to explain both ionic and covalent bonding. It’s more intuitive to understand in both cases. One can draw Lewis structures, discuss molecular shape, determine polarity for intermolecular forces, all without the orbital-based electron configurations.

(2) Do we really need hybridization? We don’t need it to draw resonance structures, and if students remember that being able to draw resonance structures contributes to stability, that’s likely good enough at the G-Chem level. Delocalization can be qualitatively discussed without hybridization through resonance structures. Many O-Chem books cover hybridization in their introduction, just enough to cover what is needed. Much of sophomore-level O-Chem doesn’t require orbitals. Geometric/steric/resonance arguments can cover a lot, and even simple nodal theory could stand in for “frontier orbital” symmetry explanations.

(3) The success of predicting the paramagnetism of O₂ comes at the cost of a whole bunch of things the students don’t understand and get seriously confused by. And frankly, the result is hardly used. Simple geometric reasoning combined with VSEPR theory can explain the anomalously weak F–F and O=O bonds. One doesn’t need paramagnetism to discuss the reactivity of O₂.

(4) Yes, the simpler model might be “wrong” but how does the orbital picture enhance student understanding of chemical reactivity? In my experience, it just adds to confusion and the students actually rely on the simpler heuristics to get them through. Again, this is at the G-Chem level. We discuss bonds breaking and forming, but viewing it through an orbital lens doesn’t confer a particular advantage.

(5) A Bohr shell/subshell model coupled with photoelectron spectroscopy can do this sufficiently without getting into the orbitals explicitly.

Now, I’d certainly want to tell students the limitations of the Bohr model (along with its strengths). Chemistry is peppered with models. We’re trying to describe the unseen. We’re forced to use models all the time. All models are wrong. Some are useful. It’s something I tell my students as we move from model to model. Goodbye, quantum numbers and orbitals. We would use Bohr electron configurations with a shell/subshell model as a foundation to bonding. So that’s roughly six hour-long class periods that could be devoted to something else.

If orbitals were not mentioned, I would still discuss atomic emission/absorption spectra and the Bohr model, and I would still discuss the strangeness of Bohr’s magic wand in the context of the Bohr model. (I could even try my Orbital Azkaban analogy!) I would only need to use a part of my “Waves and Photons” class and not worry about diffraction and interference and the quantum nature of a photon (although I would still mention the idea of a standing wave.) I could skip the photoelectric effect, wave-particle duality, the deBroglie relation, and Heisenberg’s Uncertainty Principle. I think Pauli repulsion is important, but this can be discussed without the four quantum numbers definition and talking about “space and spin” properties of fermions in simple terms. Shorter multiple bonds compared to single bonds can be attributed to Lewis-ish geometric considerations.

All that being said, I don’t expect to jettison orbitals in my G-Chem class. For one thing, our department teaches multiple sections so I might be one of ten instructors. We decide on a common textbook and coverage of topics for both semesters. I’m unlikely to win an argument to skip them especially since all of us are used to teaching G-Chem with orbitals, not to mention we also learned it that way. I also genuinely enjoy talking about it in class. I think electrons and chemical bonds are weird and wonderful – I suppose that’s why I’m a chemist. And while some students find it difficult and confusing, others enjoy delving into the “whys” – a fresh breath of air compared to the way chemistry was taught in their high school.

Just yesterday after class, a student came up to say she really enjoyed class because she experienced some gain of understanding compared to what she saw and memorized in her high school A.P. chemistry class. We had discussed ionization energy of H, He¹⁺ and He, and the effect of electron screening. We then delved into how the 1s electron differentially shields the 2s and 2p from the nucleus, and how it splits up the energy levels of orbitals with the same n quantum number. (We had previously covered the limitations of the Rydberg equation.) Then, we looked at photoelectron spectroscopy data and talked about the interplay between theory and experiment. I felt gratified for going into those details that had also excited me about chemistry when I was an undergraduate!

I leave you with a picture from one of my opening slides from this past week: the orbitron gallery. Check out the WebElements store if you want to get your own!