Teaching about chemical bonding

Chemical bonding is an area where a good many student alternative conceptions have been identified, and some of these tend to be both very common and tenacious

(that is, they tend to persist even when teachers point out they are scientifically incorrect). Moreover, these are not ideas that can develop from direct experience - as happens with the very common alternative conception that a moving object will only have a limited impetus and so will spontaneously stop moving sooner or later. Physics teachers struggle to teach the formal scientific model in that case because the ‘misconception’ is actually a very reasonable generalisation of what we all seem to experience in the real world, where we cannot exclude friction and air resonance and other forces which operate on a moving object without being obvious to an observer. (In terms of Figure 3.1, there is a ‘grounded learning impediment’). Yet children do not come to class with a wealth of experience of observations of chemical bonds. Chemical bonding is a concept at the theoretical, submicroscopic level (see Chapter 5), which has to be related to observable properties of actual samples of substances.

Nor, generally, are alternative conceptions about molecules and bonds due directly to ideas that have high currency in students’ social worlds that they then bring to the chemistry class. That is, whereas the idea that exercise can give you energy clearly is part of an everyday discourse, students do not get told that chemical bonds form so that atoms can have full electron shells in the home or playground: such ideas must develop in the process of learning the chemistry itself. Students may develop ideas, inter alia (Taber, 2012a):

■ that hydrogen bonds are a class of covalent bond;

■ that in ionic structures a singly charged ion (e.g. Na+) can only form an ionic bond with one counter ion (e.g., Cl-);

■ that dative bonds are covalent bonds (rather than something between a pure covalent and pure ionic bond);

■ that there is no actual chemical bonding in metals;

■ that on homolytic bond fission, each atom always gets “its own” electron back

A critic of chemistry teaching would perhaps conclude that these ideas can only have derived from teaching, and - therefore - if students so commonly develop alternative conceptions of the chemistry, then chemistry teachers have not been doing a very good job.

Sometimes, perhaps, teachers themselves hold alternative conceptions of some of the science and so teach incorrect ideas in their classes (Taber & Tan, 2011). However, what seems more often to happen is that students interpret teaching in a different way to that intended. According to the constructivist perspective of how learning occurs, students have to make sense of teaching - including teaching about theoretical objects and abstract ideas like molecules and bonding - in terms of the ideas and experiences that are already familiar (Taber, 2014). Bonds are often understood in terms of more familiar material links - glue, rivets, chains, elastic bands, cocktail sticks - and the driving force for bonding is often understood in anthropomorphic terms: what atoms or electrons want or need or prefer! Teachers should therefore think about how students will understand teaching and should be constantly checking how students are making sense of abstract ideas in lessons (e.g., see Taber, 2018a) - and seeking to carefully shift their thinking towards more scientific accounts.

A principled approach that supports teaching

From a scientific point of view, bonding can largely be understood in term of the forces between charged particles, and how these tend to bring about more stable (lower energy) configurations. Chemical species that tend to form and stay around are equilibrium arrangements where the charges are configured in a way that they cannot easily be nudged into other more stable arrangements. Such arrangements can be disrupted by inputs of sufficient energy, but once that energy is dissipated, the particles return to stable arrangements.

The charged particles concerned could be considered protons and electrons - but (in chemistry) we can usually simplify our thinking about these systems. This is just as well, as teaching about the structure of the methane molecule as an arrangement of 16 protons and 16 electrons would clearly be challenging for most students. Protons are found in nuclei, and although nuclear processes (fission, fusion, radioactivity) are important, they are beyond introductory chemistry (and some would argue actually outside of chemistry completely). So, our methane molecule simplifies somewhat to five nuclei and ten electrons, and a tetrachlo- romethane molecule to five nuclei and 74 electrons. Yet, often in chemistry, and indeed nearly always in introductory chemistry, we can assume that inner shell electrons remain largely unaffected by chemical change. We can consider chemical systems to be composed of atomic cores (nuclei and inner shell electrons) and the valence electrons from the outer shells associated with those atomic cores. Now our methane molecule is simplified to five atomic cores and eight electrons, and our tetrachloromethane molecule to five atomic cores and 32 electrons (see Figure 10.4).

Once again, we should keep in mind that we are dealing with models, and models are simplifications, and that simplifications have limited ranges of application before their usefulness breaks down. So, it is not always possible to consider atomic cores as discrete and impervious components of chemical systems. For example, in understanding the bonding in transition metals, there is more going on than metallic bonding involving the outer shell valence electrons being delocalised within an array of atomic cores (i.e., metallic cations). Electrons from the ‘core’ are involved in some degree of bonding interactions as well, explaining some of the properties of these elements (such as high melting temperature). That is a complication that need not be discussed when introducing bonding ideas, although it is helpful if students appreciate that they are being taught a model, a

One possible way of representing molecular structures (cf. Figures 4.1,6.1)

Figure 10.4 One possible way of representing molecular structures (cf. Figures 4.1,6.1): any choice of representation simplifies; emphasises (and de-emphasises) particular features; and reflects a partial model of the structures represented

simplification that is useful at times but may need to later be developed for some purposes.

Often textbook figures of molecules of compounds use different symbols (e.g., • and X) to show the electrons ‘from’ the different atoms, but students often infer this implies some kind of ownership, and therefore perpetual association of the electrons with their ‘own’ atoms - so they would be retained on bond fission - even though the atoms no longer exist in the molecular structure. Indeed, the common practice (as in Figure 10.4, cf. Figure 10.5) to show molecules as formed by overlap of atoms may itself mislead learners to think of the molecule as a temporary association of continuing atoms. Any representation is potentially going to mislead students (as a ‘realistic’ representation of a molecule is not possible, see Chapter 5) unless such issues are explored.

The quantised elephant in the room

There is a more significant omission in teaching a model based on forces between charges, and that is that there are the quantum rules that set out restrictions on the configurations that the quanticles can take up because variables such as energy and angular momentum are quantised and therefore only found in multiples of the basic quantum. This is a very important aspect of the basic nature of matter. It is not usually formally taught until late in secondary school chemistry (in elective courses) and tends to be considered too abstract and difficult for most school-age learners.

Electrons are found in shells, with a maximum of two electrons in the first shell, eight in the second, eighteen in the third, and so forth. Moreover, introductory

A slightly different representation to Figure 10.4, perhaps better suggesting that the atoms as such no longer exist once a molecule is formed

Figure 10.5 A slightly different representation to Figure 10.4, perhaps better suggesting that the atoms as such no longer exist once a molecule is formed.

chemistry tends to focus on the elements in period 1-3, so that every compound met involves elements which, other than hydrogen, have atoms which seem to ‘fill-up’ with eight electrons in the outer shell. This may seem to be the case in period 3 because younger students are not usually introduced to examples like PC15 or IF7, or asked to think about the structure and bonding in the sulphate ion, S042-, in their introductory chemistry courses.

Anthropomorphic misconceptions

In the absence of any scientific explanation for the ubiquity of octet structures, students tend to anthropomorphise: atoms want full shells, they need full shells, they desire full shells, they will (intentionally!) get involved in chemical processes to acquire full shells (Taber & Watts, 1996). The danger of such ways of thinking is that they stand in the place of genuine scientific explanations.

Students asked to explain reactions such as

or

will tend to explain these reactions occur because hydrogen needs to have 2 electrons in its outer shell and the other elements want to have full outer shells of eight electrons. Usually students are so convinced that this is a sound explanation that they completely ignore how the reactant molecules ‘obey’ this octet rule just as much as the product molecules, and so the same logic could just as easily ‘explain’ the (non-feasible) reverse reactions:

or

What exactly does it mean to ‘share’your electrons?

Often, teaching schemes begin with covalent bonding and use the metaphor of a covalent bond being a ‘shared’ pair of electrons - sometimes without exploring what exactly sharing means in this context. A teacher cannot be surprised that students may later talk of atoms ‘stealing’ electrons and the like if a social metaphor is presented as if a satisfactory description. As suggested earlier, it would be better to focus on how the negative electrons can hold the positively charged atomic cores in the structure when there is a balance of attractive and repulsive forces. A student might reasonably ask how the two negative electrons can be considered to behave as a pair if they repel each other - a question that cannot be addressed in any detail without invoking that elephant we may prefer to pretend is not in the room.

Covalent bonding is often illustrated with misleading hypothetical schemes for how the molecules came about: thus, methane is often shown as being formed from an isolated carbon atom and four isolated hydrogen atoms - allowing students to acquire the mistaken notion that bonds form in order to allow atoms to follow the octet rule (chemical reaction mechanisms are discussed in Chapter 11). The students may be new enough to the subject not to spot this trick and so to ask where these radical atoms originated: a teacher should know better than to use such deceptive and unscientific devices in explanations. A teacher should only explain a reaction as starting with non-bonded atoms if they can show the students that reaction using reactants in that form - but I doubt any school chemical stores can supply such reagents!

The ionic bond is based on the attractions between charges

The traditional teaching scheme will often then move on to ionic bonding. This is a more complex situation because it does not involve discrete links between adjacent atomic cores but rather relates to the overall regular structures of myriad cations and anions that allow forces to be balanced by placing positive ions closer to negative ions than other positive ions and vice versa. (That may seem to give an overall attractive force, not an equilibrium, but one has to consider that once ions come very close, then the outer electrons, all negative, have a small separation so contribute strong repulsions.) Often, to simplify matters, teaching models ignore the lattice originally and focus on molecule-like parts of the structure - such as one

Na+ and one Cl" ion in NaCl. Whilst this is certainly simpler, the aim is to explain NaCl, which has an extensive lattice of ions, not hypothetical Na+-Cl“ ion pairs (which may exist at a low level in the vapour phase, or hydrated as a minority component in very concentrated solutions - but have no role in the solid). If we can treat water as composed of H20 molecules to a good first approximation (e.g., ignoring the very low proportion of ions present, see Chapter 8), then there is no place for Na+-Cl" ion pairs in any introductory model of sodium chloride.

Ionic bonding is not about electron transfer

A common scheme imagines that these irrelevant Na+-Cl~ ion pairs form by electron transfer from a single isolated sodium atom to a single isolated chlorine atom. There seems to be something in our psychology which wants to explain origins from some kind of perfect starting point (here, the fallacy of initial atomicity) as if elements were originally created or formed as isolated atoms (they were not) which then entered into bonding arrangements. The immediate problems with this scheme are that it both refers to a chemically non-viable set of reagents (neither sodium nor chlorine is available in the laboratory in an atomised form: rather sodium is metal and chlorine is molecular-) and that it is energetically non-viable. It takes more energy to remove an electron from a sodium atom (its first ionisation enthalpy) than would be recovered by forming a chloride ion (its electron affinity) - so if one had a Na+-Cl~ ion pair that was not stabilised (e.g., by being solvated) it would spontaneously break down to give the atoms.

The common consequence of being taught this scheme is that students often think that

■ there are molecule-like ion pairs in NaCl;

■ each ion can only form one ionic bond (produced through one electron transfer event);

■ there are two types of bond in the structure (ionic bonds within ion pairs and just forces between ion pairs);

■ the ion pairs exist in the molten or solvated NaCl.

If this model were true, then NaCl should have a low melting temperature and should not be an electrolyte! As our submicroscopic models are used to explain macroscopic features (see Chapter 5), there is little value in a false model which leads us to infer the wrong properties of NaCl.

Preparing NaCl (with no need for electron transfer)

Students can prepare NaCl in the laboratory by neutralisation, followed by evaporation of the solvent. The product, NaCl, contains Na+ ions which were present in the NaOH reagent solution and СГ ions that were generated from the HC1 solution. Despite this, it is very common for students (even if they have undertaken this preparation) to claim that the ions in NaCl were formed when sodium atoms gave their electrons to chlorine atoms. This seems to be an impression given by some textbooks or even teachers.

This is a very odd state of affairs. A fantastic account of ions being produced from non-existent atoms is certainly not a simplification of the chemistry - we can explain the ions as simply being present in the bottles of the reagents we used in the preparation! The scientific explanation is that the ions that were already present in the mixture were attracted together as the solvent evaporated. It seems genuinely mischievous to instead teach a more complex account that we know is false.

A balance of attractive and repulsive forces - a common core for bonding

The other problem with using both the sharing metaphor and the false ‘electron transfer’ notion as core ways of explaining these bonding types is that this obscures the basic commonality that all forms of chemical bonding can be understood as charged particles coming together until there is a balance of attractions and repulsions. Metallic bonding may be taught as involving a ‘sea’ of electrons. This is not a poor metaphor in itself but may become seen as an explanation rather than a description. If all bonding is seen as either electron sharing or electron transfer, metallic bonding tends to be understood in terms of these patterns (Taber, 2003c).

Moreover, as covalent and ionic bonding seem quite different in nature, the distinction is treated as a dichotomy rather than a continuum, and dative and other polar bonds are not later readily understood as on a spectrum between pure covalent bonds and (ideal) pure ionic bonds. Other common consequences are that hydrogen bonds are suspected of being covalent bonds; and weaker interactions such as van der Waals’ forces are not readily understood as being a kind of bonding.

The ‘sea of electrons’ metaphor offers a useful image to help student visualise an unfamiliar and abstract structure at a scale they cannot observe. However, it should be treated as a familiarisation device to help introduce an abstract idea (Taber, 2018a), and it should be made clear to students that it is not an explanation of metallic bonding - just as ‘sharing’ electrons does not explain the covalent bond. A teacher needs to make these points explicitly to students, who will often not appreciate them spontaneously.

Thinking about a teaching order

There is a strong argument for saying it would be better to teach bonding starting with the simplest cases - elements with only one kind of bonding to consider (Taber, 2001). We might first teach about solid argon (weak bonding between single atoms), pure metals (arrays of cores with delocalised electrons), and giant molecular structures such as diamond. Starting with an example like solid argon helps learners focus on the electromagnetic nature of the chemical bond, making it more difficult for students to develop the common notion that chemical bonding is about acquiring full electron shells.

Graphite is more complex, as we have to invoke a second kind of bonding (it has the kind of bonding in diamond but also the kind in solid argon). This is similar to bonding found in sulphur (S8) or phosphorus (e.g., P4) or in solid oxygen or nitrogen. Then bonding in compounds might be introduced, such as in silica (similar to diamond) or in methane or in ionic salts. A sequence such as this can help the teacher emphasise the fundamental principle common to all forms of bonding.

 
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