Misunderstanding the ionic bond

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Taber, K. S. (1994) Misunderstanding the ionic bond, Education in Chemistry, 31 (4), pp.100-103.

Although Education in Chemistry has a web-site which includes recent articles form the magazine, this does not go back as far as 1994. Content of older issues of Education in Chemistry has been scanned and included in the Royal Society of Chemistry's historical on-line archive, available to members, but not freely available to the general public. (For those with access, the pdf can be downloaded from https://hc-content.rsc.org/EIC/EIC-1994-31-4-03/EIC-1994-31-4-03.pdf.)

The text of the article follows:

Students beginning A-level chemistry seem to have misconceptions about the nature of ionic bonding. Could there be implications for the way in which bonding is taught at Key Stage 4 and A-level?

At the beginning of their A-level chemistry course, I invited some of my students to take part in a study of how their understanding of chemical bonding developed during the course. The 10 'volunteers' were inter- viewed during the first few weeks of term to find out how much they knew and under- stood about chemical bonding at the start of the course. As might be expected, there were variations in the breadth of knowledge and depth of understanding. One aspect of particular interest was that, although these students were able to discuss ionic bonding, many of their comments implied that their mental models were incorrect. Moreover, there seemed to be similarities between the wrong ideas revealed – even though the students concerned came to Havering College from different schools. If these common misconceptions are more than just coincidence, and are found to be wide- spread, then there could be implications for the teaching of chemistry in schools, colleges and sixth forms.

Alternative conceptions

There has been a great deal of research into children's ideas and understanding in a number of scientific topic areas. For example, the Children's Learning in Science Project (CLiSP) has produced materials about understanding particle theory ¹ and basic chemistry.² The viewpoint underlying such research is that learners often construct their own meanings and understandings, based as much on what they experience outside the classroom as within – but much of this 'knowledge' is actually at variance with accepted science, and can impede orthodox understanding.

The learner's own ideas may be fragmentary and only weakly held to, or they may comprise an extensive and coherent frame- work of concepts that has much explanatory value for the learner. Researchers have used a variety of terms to describe such alternative ideas – such as misconceptions, alternative conceptions, children's science, alternative frameworks, intuitive theories, informal ideas, theories in action, and mini theories. Despite a large literature covering many aspects of school science there has been little attention focused on the topic of bonding. ³

[Focal figure used in interviews with A level (an elective course typically taken by 16-19 year old learners in England).]

I have a particular interest in studying students' understanding of chemical bonding and its development during an A-level course. The main research technique is a sequence of in-depth interviews, focused on figures showing atoms, molecules, lattices etc. Most of the comments discussed in this article relate to Fig. 1, which represents a layer of ions in NaCl. ⁴ The students are considered to be colearners with the teacher, both learning from the interview/feedback. Students are then re-interviewed as the course progresses.

By analysing the interviews we can build up a case study of how a particular colearner develops an understanding of chemical bonding. ⁵ It is taken as axiomatic that each learner's network of ideas is unique, and case studies are not intended to be generalised to the wider population. However, where research evidence suggests that a large number of different learners could share similar alternative constructions, it may be useful to model the common elements, which may help other teachers to diagnose misconceptions, and could possibly suggest how to avoid them. For example, one researcher was able to suggest a series of common alternative frameworks that young people exhibited about force ⁶ and energy. ⁷

This article concentrates on the comments of five of the colearners at the start of their course. These students (signified by the letters K, N, P, Q and T) appeared to exhibit similar 'misconceptions' about ionic bonding. The relevant parts of the taped interviews were transcribed (see Box I for an example), and then vignettes were edited together from the colearners' comments. For example, Box 2 presents an abridged vignette of P's comments when discussing Fig. 1.

Ionic bonding is …

The orthodox framework for understanding ionic bonding at GCSE and A-level may be described simply enough. The ionic lattice is bound by the attractions between oppositely charged ions. The attraction between anion and cation is the bond, and its strength depends on the size of the ionic charges and their separation. The ions are atoms or molecules that have an excess or deficit of electrons, and the origin of this imbalance is not important in understanding the forces involved in the bonding. I will call this simple model the 'electrostatic framework'.

The common threads of the alternative framework – that is the one that students often use – may be summarised in three conjectures that students seem to make.

The valency conjecture: the atomic electronic configuration determines the number of ionic bonds formed. For example, a sodium atom can only donate one electron, so it can only form an ionic bond to one chlorine atom.
The history conjecture: bonds are only formed between atoms that donate/accept electrons. For example, in sodium chloride a chloride ion is bonded to the specific sodium that donated an electron to that particular anion, and vice versa.
The just forces conjecture: ions interact with the counterions around them, but for those not ionically bonded these interactions are just forces. For example, in sodium chloride, a chloride ion is bonded to one sodium ion, and attracted to a further five sodium ions, but just by forces – not bonds.

I call these conjectures the 'molecular framework', because it appears that colearners perceive small groups of ions within an ionic lattice to be bonded to one another, but only attracted to surrounding units – as in a molecular solid, such as iodine or sulphur. It is important to emphasise that each learner will have a unique conceptual structure, and the molecular framework is only a model to highlight the similarities. Although a co- learner may explicitly refer to molecules Within an ionic lattice, this was not considered to be a major criterion for this framework – rather , it is the implicit molecular nature of the valency, history and just forces conjectures that is considered to characterise this model. Research has also shown that it is not uncommon for learners to have several conflicting frameworks, and several of the colearners interviewed appeared to be in transition (or confusion) between the electrostatic and molecular frameworks.

'I' represents me, the interviewer.

I: Can you identify any bonds in this diagram; can you point to where the bonds are?

P: [Pause, ca 4 s.] Probably. [Pause, ca 4 s.] Maybe these two are attracted, like this. So, one would be there, then, suppose there was one between there and there, there.

I: Er, right, let me, let me pick on an ion. This sodium ion here which is fairly near the middle. How many chlorides is it bonded to?

P: One.

I: Do you know which one it is?

P: Then how come it's got four? [Pause, ca 3s.]

I: Do you think the diagram is maybe inaccurate? You know , I drew these diagrams, so they might not be very accurate.

P: Yes I think so, because how is it possible, if this is like saying, can we just take, that bit?

I: Mmm.

P: There, it's like saying, Na's in the middle.

I: Mmm.

P: And you've got four Cls there.

I: Mmm.

P: But how can you have that when there's only one electron in the sodium, to give to one chlorine, so how are these three – how are they attached to that?
[Content of] Box 1: Extract of interview transcript

When asked how many chlorines were bonded to a sodium , P replied , 'one', but then added , 'then how, how come it's got four?' So did she think the diagram was inaccurate? 'Yes , I think so, because how is it possible…it's like saying, Na's in the middle, and you've got four Cls there, but how can you have that when, there's only one electron in the sodium, to give to one chlorine, so how are these three, how are they attached to that?'. A chloride would be bonded to 'one' sodium, although she couldn't tell which one it was in the diagram. If we did know which Na had donated an electron to which Cl, then we could locate a definite bond where those ions touched. For adjacent ions that had not been involved in ion transfer 'it's possible' 'they could have like a weak attraction', 'because that's positive and that's negative'. However, 'it's a possibility that those would be, stronger…actually where the electron has been transferred '. The other interactions would not be called ionic bonds, but were 'just a weak attraction' due to 'the positive charge on the sodium, and the negative charge on the chlorine, 'because as we know, positive and negative attract'.

Box 2: Colearner's comments about ionic bonding

Some discussions

P exhibited all three main aspect of the molecular framework, and was confident enough in her mental model to doubt that Fig. I was accurate. K also appeared to use the valency, history and just forces conjectures, and explained his perception of the inaccuracy of Fig. I as a distortion arising from representing a three dimensional structure in two dimensions. In reality, he believed each ion only had one nearest neighbour. N starts to answer from the alternative 'molecular' interpretation (seeming to use both the history and just forces conjectures), but then changed to the more conventional 'electrostatic' framework as she thought through her answers. Q demonstrated a variation, in that although a valency conjecture was made, it was not applied to Fig. I because she judged that no bonds were present – just forces. For Q, electronegativity and charge are both somehow involved as causes of the lattice forces, and the importance of electron transfer is emphasised. It is interesting that Q is able to talk about the bonding being ionic, and yet she can still deny that a sodium ion in Fig. I is bonded to anything. T shows all the main feature of the 'molecular' framework for understanding ionic bonding – even making explicit use of the term 'molecule'. He seems uneasy about the 'just forces' conjecture, and has equal forces (though not equal bonding) between all adjacent ions, but he has not yet learnt to operate from an 'electrostatic viewpoint'.

It has not yet been established how widespread the alternative molecular framework for understanding ionic bonding is. If many young people's ideas are closer to the molecular framework than the orthodox electrostatic framework, then it is important to understand why. I suggest that the following three factors may encourage young people to adopt the molecular framework.

First, the standard presentation of ionic bonding tends to have three stages: (i) drawings of the electronic structures of single atoms of sodium and chlorine; (ii) the transfer of an electron from sodium to chlorine; and (iii) the electronic structures of the ion pair. This approach emphasises the electron transfer and implies a special relationship between the two ions in the ion pair. In fact, this is a pseudo-molecular presentation, and in some books the lattice structure formed is completely ignored. It also avoids the fact that even when sodium chloride is formed by binary synthesis, the chlorine does not exist as isolated atoms before the reaction, and the electron that chlorine accept are donated by the metallic conduction band, not individual atoms.

Indeed, the 'true history' of many ionic substances would show them to be formed by evaporating a solution, where the cations have donated electrons to, and the anions have accepted electrons from, completely different species from those in the compound being considered. Although authors and teachers do not intend to imply a molecular nature, it may come across as such. The recent debate in Education in Chemistry has shown that some chemistry teachers do feel that they need to have a term for the ionic 'molecule' within the lattice. ^{9 12}

My second suggestion is that, in explaining ionic bonding, the teacher will be making use of his or her tacit knowledge about electrostatics – about charge, force and energy. Although the learner will also have meanings for these terms, they may be very different from those of the teacher – research has shown that alternative conceptions abound in these areas. For example, the teacher may visualise electron transfer between sodium and chlorine as follows. The sodium valence electron is influenced by the greater core charge on the chlorine compared with the sodium, giving rise to a resultant force that causes the transfer.

For the pupils, however, the cause of electron transfer may be seen as a tendency (or 'need') for atoms to attain a noble gas configuration. Although this tendency has an electrostatic basis, it is usually presented simply as: 'full outer shells are stable'.

A third factor could be the order in which ideas are presented. Bonding is a highly abstract topic for most youngsters at KS4. The learner is often first taught about covalent bonding, with its emphasis on valency and molecules. If, after an effort to make sense of this material, the learner is then asked to understand ionic bonding, it is perhaps not surprising if covalent bonding is used as a basis for understanding this new knowledge. Molecular ideas may be adopted by analogy with the covalent case.

In a recent article in Education in Chemistry, Michael Laing argued that not enough is done to distinguish between macromolecular (giant molecular) structures, and simple molecular solids, where the lattice integrity arises from van der Waals forces. ¹³ Perhaps the three factors mentioned could be further compounded if pupils have a single mental model of covalent solids – i.e., some form of amalgam of the discrete molecules in iodine, and the extended covalent network in diamond.

Conclusion

My research has shown that successful GCSE students may begin A-level chemistry with significant misconceptions of ionic bonding. The alternative framework presented is based on in-depth interviews with a small number of students. Further work is needed to establish how common such ideas are, and to investigate the suggested connection with the 'pseudo-molecular' presentation often used in teaching.

In the meantime, teachers may want to consider the following suggestions about introducing ionic bonding.

Focus on the electrostatic lattice forces, rather than ion formation.
Clearly distinguish between ion formation (electron transfer), and ionic bonding.
Do not restrict diagrams to ones showing molecule-like entities (pairs of atoms and pairs of ions), but include ensembles of ions.
If the 'reason' given for ion formation is the stability of noble gas electronic configurations, then make sure that this is not also considered sufficient reason for the subsequent formation of bonds between ions.
Discuss the differences (as well as similarities) between lattices held together by ionic, covalent and intermolecular forces.
Include an example of an ionic material formed via precipitation, e.g., barium sulphate, to emphasise that ionic bonds can form even if no electron transfer takes place between the barium and sulphate species.
If the term valency is used at all, then discuss explicitly the meaning of electrovalency in terms of ionic charge formed, and compare this with covalency; and make it clear that the number of ionic bonds formed is not determined by electrovalency.

Those teaching at A-level and beyond may also want to consider this list when 'recapping' basic ideas before introducing new bonding concepts such as electronegativity.

The general conclusions that I would like to draw are familiar to most teachers. First, to remember that students do not share the extended network of background we call upon when teaching topics (e.g., electrostatics in teaching bonding). Secondly, that it is unwise to make assumptions about prior knowledge just because our students seem intelligent and motivated, and have passed earlier examinations. Finally, I suggest that the best way of investigating our students' ideas is to talk to them – but only if we are prepared to be surprised by some of their comments!

The material in this article formed part of an Education Division lecture at the 1993 Autumn Meeting of the RSC.

Keith S. Taber [was] a senior lecturer at Havering College of Further and Higher Education…

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