Avoiding bonding misconceptions

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Students’ understanding regresses after teachers introduce the octet rule

Bonding is one of the most challenging topics for chemistry students. The so-called ‘octet rule’ can cause problems. Often teaching and understanding of bonding is based on the notion of atoms ‘needing’ full outer shells, with the octet rule used as an explanatory principle rather than a rule-of-thumb.

Many textbooks describe the octet rule as an explanatory principle, and even exam board specifications have included content phrased incorrectly. Coupled with the fact teachers are likely to encounter erroneous use of the octet rule during their training, it is no wonder many (including myself, on reflection) lead students to develop faulty conceptual understanding.

Many textbooks describe the octet rule as an explanatory principle, and even exam board specifications have included content phrased incorrectly.¹ Coupled with the fact teachers are likely to encounter erroneous use of the octet rule during their training, it is no wonder many (including myself, on reflection) lead students to develop faulty conceptual understanding.

Jarkko Joki and Maija Aksela have followed a student cohort’s progress in understanding bonding throughout their secondary education. In 2015, the pair from the University of Helsinki in Finland interviewed students after lower-secondary level. The cohort’s teachers had focused on electrostatic interactions as opposed to the octet framework, and these students generally had a well-constructed conceptual understanding of bonding based on electrostatics.

In their latest study, the researchers interviewed students again after upper-secondary school. They investigated the students’ resistance to the ‘octet framework’ as a result of their earlier exposure to more robust explanations. They interviewed eight students, now at three different schools along with three of their five upper-secondary chemistry teachers. While the authors acknowledge the limitations of such a small sample, the interview excerpts provide fascinating insight into the thought processes of individuals as they address challenging questions.

In their latest study,² the researchers interviewed students again after upper-secondary school. They investigated the students’ resistance to the ‘octet framework’ as a result of their earlier exposure to more robust explanations. They interviewed eight students, now at three different schools along with three of their five upper-secondary chemistry teachers. While the authors acknowledge the limitations of such a small sample, the interview excerpts provide fascinating insight into the thought processes of individuals as they address challenging questions.

The interviews found the students’ apparent understanding of bonding had regressed since lower-secondary level. Most had adopted an octet framework linked to a minimum energy principle to explain bonding. The students involved were strong students expected to get good grades.

Interviews with teachers showed they used the octet rule as an explanatory principle without detailing how it relates to an energetically favourable situation or to the electrostatic basis of bonding. Even where a teacher recognised the limitations of the octet rule, they did not feel compelled to discuss this with students unless specifically asked about it. The student outcomes can be linked to the impact of the teaching they received.

Teaching tips

Ultimately, we all need to unpick our understanding of bonding. We should consider the role of electrostatics and energetics in determining the bonding in a substance, and how this relates to the fact bonded atoms do generally have full outer shells.

Introduce all forms of bonding as electrostatic phenomena. This is straightforward with ionic bonding, and metallic bonding to an extent, but covalent bonding has traditionally been blighted by the ‘it’s a shared pair of electrons’ definition.
Tell students the octet rule is a useful rule-of-thumb or memory-aid for working out ion charges and formulae of compounds, but that it should never be used as an explanation.
Remember that when an atom such as fluorine gains an electron, this is attracted into a low energy orbital by the protons in the nucleus, while an electron being lost from an atom such as sodium experiences a weaker force of attraction than those in the complete inner shells due to distance and shielding. The resulting ions attract each other and form an ionic lattice. These processes are not driven by a desire for full shells, but by energetics.
Lessons on Born–Haber cycles are a good time to get students to think about why compounds with ions such as Na²⁺ are unfeasible on energetic grounds, ie the energy required to remove a second electron far outweighs any gain in lattice enthalpy. You can also show formation of compounds containing F^2- are unfeasible, which helps illustrate why we generally observe full outer shells in atoms or ions in compounds.

David Read