Chemical Bonding Confusion

The current issue of Education in Chemistry has a great article on the issues surrounding teaching chemical bonding with some useful links to additional resources on Learn Chemistry. Several misconceptions stem from the dominance of the octet/filled shell model in the description of chemical bonding, where the over-riding consideration is the “need” to have a full outer shell.* With only this consideration in mind, Stephenson and Warren explain in their article, students can easily come to the conclusion that reactions happen due to the transfer of electrons to form ionic bonds or share electrons to form covalent bonds.

This isn’t a new phenomenon. In Education in Chemistry in 1996, Ronald Gillespie wrote that relying solely on orbitals to explain chemical bonding:

in introductory courses is fraught with many difficulties and at times may even obscure the fundamental reason for the chemical bond—the electrostatic attraction between positive nuclei and negative electrons.

Nor is this an issue restricted to school level chemistry. There is a lot of research demonstrating significant misunderstanding at third level. An interesting paper in Chemistry Education Research and Practice last year aimed to investigate upper secondary and early university students’ understanding of bonding by asking them to develop their own models after they had been taught about ionic and covalent bonding.

In this research study, nine out of twenty-four students defined bonding in terms of “giving up” (ionic) or “sharing” (covalent) electrons. Students diligently reproduced models that we would commonly see in text books (and oh, I confess, our lecture notes). But in discussing these models, the emphasis was on the filling of shells as being responsible for the bond, rather than the electrostatic attraction to between electrons and nuclei. An additional component to this observation was the lack of use of electronegativity to determine whether ionic or covalent bonding is occurring.

An alternative approach?

In an article on this topic in Education in Chemistry in 2011—required reading really—Keith Taber describes the use of the electrostatic model in explaining many interactions in chemistry, not least that between a nucleus and electrons.  When considering covalent bonding, for example, each atom in the bond will have electron density around its positive core. The attraction between the cores and the electrons balances out the repulsion between the cores. (The activity “Why do hydrogen and fluorine react?” linked in Stephenson and Warrens’ article may be a good activity to develop this concept with students).  A similar conceptualisation of an ionic lattice as one where cations and anions are interspersed and arrange close enough so that the attractive and repulsive forces balance out. The article considers many other phenomena that can be considered from an electrostatic model.

Confusion reigns

As I write this post, I’ve just consulted a textbook common at introductory college level. The chapter on chemical bonding opens with a cartoon on the different types of bonding, with ionic having the label “electrostatic attraction” and covalent having the label “electrons shared”. The main text for ionic bonding includes the terms: “reactions of non-metals and metals”, “electron transfer”, “low ionization energy”, “high electron affinity”, “valence electron” along with electron configurations and both a single +/— ion pair as well as a crystalline lattice. This is all in the space of about half a page! Covalent bonding is then presented, with a clear diagram showing the push/pull of attraction and repulsion. Only then are electronegativity and polar-covalent bonding introduced, and then finally, a section, with no diagrams, on the fact that the nature of bonding is more correctly considered as a continuum.

I’d be interesting in hearing other educator’s views on whether and how the introduction of these concepts should be reversed.

Note: *A separate topic, but an interesting article on considering atomic “desires” is summarised by David Read in a recent EiC article – see link.