Chemical bonding is one of the key concept areas in the subject, and is also an area where learners are known to commonly acquire alternative conceptions
Some of these alternative ideas are considered in this chapter, with suggestions for improving the teaching of the topic.
The full shells explanatory principle
Students are found to commonly use the octet rule - a useful heuristic for identifying stable chemical species - as the basis of a principle to explain chemical reactions and chemical bonding.’
According to this ’full shells explanatory principle’ bonding occurs ‘in order [for atoms] to try to achieve a stable structure ie 8 electrons in the outer shell of the atom’.
Students relate the ’sharing’ of electrons in covalent bonds to the full shells explanatory principle, so that ’the electrons are shared to create a full outer shell’, and the ’covalent bond is the sharing of electrons to complete full valency shells’.
The full shells explanatory principle is inherently anthropomorphic, as no physical force is invoked to explain why systems should evolve toward certain electronic configurations. Rather, it is assumed that this was what atoms ‘want’, and so they act accordingly.
Given that the starting point for many students’ thinking about bonding is the atoms’ perceived need to achieve a full shell, it is not surprising that often
- Students see chemical bonding and forces within chemical structures as largely unrelated.
- Students limit their category of chemical bond to types of interactions that can be readily conceptualised in terms of the full shells explanatory principle.
When is a chemical bond not a force?
We could define a chemical bond as that which holds the parts of a chemical structure together.
It is sensible to start this chapter by exploring what a chemical bond is, I would suggest it is a force which holds chemical species together. This force can usually be considered as an electrical interaction. The source of this interaction is the nature of chemical species themselves - composed of positively charged nuclei and negatively charged electrons.
It will be noted that this description would appear to be just as applicable to individual atoms or ions as to molecules and lattices.
If there were no quantum restrictions on where electrons could be located in chemical systems, then teaching about and studying chemical bonding would be much simpler - we would not have different categories of chemical bond such as covalent and hydrogen bonding.
Students’ multiple conceptual frameworks for bonding
Just as teachers will use multiple models of bonding to help learners appreciate the abstract ideas involved, so students may develop manifold conceptions of chemical bonds. At some point, successful post-1 6 students are able to move beyond notions of bonds as shared electrons, to see bonds as electrical interactions.
Research suggests that the alternative ideas sometimes co-exist alongside developing more sophisticated ~understandings.~ For example, it is often found that post-I 6 students are in transition between two models of the ionic bond.
Student definitions of bonds: the bonding dichotomy
Students seem to acquire this dichotomous classification of bonds readily, and when they do it means that they do not see bonding as primarily an electrical phenomena. Once this scheme has become established the student finds it difficult to appreciate bonding that is intermediate (polar bonds) or falls outside (eg hydrogen bonding) this narrow definition of bonding. Now clearly covalent and ionic bonds are very significant bond types, as many important substances can be understood to have - to a first approximation - either ionic or covalent bonding. However, the effect of pupils in school learning about bonding as a dichotomy of these two types, is to act as an impediment to later learning.
|Electrons are shared between non-metal atoms||Electrons are transferred from metal to non-metal atoms|
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These PDFs have been taken from the popular book, Chemical Misconceptions : Prevention, diagnosis and care: Theoretical background, Volume 1, by Keith Taber.
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