Category: octet rule framework

Chlorine atoms share electrons to fill in their shells

Umar was a participant in the Understanding Chemical Bonding project. When I spoke to him in the first term of his course he was unsure whether tetrachloromethane (CCl4) would have ionic or covalent bonding.

When I spoke to him near the start of his second term, I asked him again about this. Umar then thought this compound would have polar bonding, however he seemed to have difficulty explaining what this meant ⚗︎ . Given his apparently confused notion about the C-Cl bond I decided to turn the conversation to a covalent bond which I knew, well certainly believed, was more familiar to him.

Is it possible for chlorine to form a bond with another chlorine?

[Pause, c.2s]

Yeah.

What substance would you get if two chlorine atoms formed a bond?

[Pause, c.2s]

You get, it still, you get, if you had like two chlorines it depends what groups are attached to it, to see how electronegative or electropositive they are.

What about if you just had two chlorine atoms joined together and nothing else, is that possible?

[Pause, c.3s]

No.

No?

On their own.

Not on their own?

No.

Umar's response here rather surprised me, as I was pretty confident that Umar had met chlorine as an element, and would know it was comprised of diatomic molecules: Cl2.

So you couldn’t have sort of Cl2, a molecule of Cl2?

[Pause, c.1s]

Yeah, you could do.

Could you?

[Pause, c.2s]

They might be just, they might be like, be covalently bonded.

Perhaps the earlier context of talking about polar bonds and the trichloroethane molecule somehow acted as a kind of impediment to Umar remembering about the chlorine molecule. It seemed that my explicit reference to the formula, Cl2, (eventually) activated his knowledge of the molecule bringing to mind something he had forgotten. Although he suggested the bond was (actually "might be") covalent, this seemed less something that he confidently recalled, than something he was inferring from what he could remember – or perhaps even guessing at what seemed reasonable: "they might be just, they might be like, be covalently bonded".

As often happens in talking to learners in depth about their ideas it becomes clear that thinking of students 'knowing' or 'not knowing' particular things is a fairly inadequate way of conceptualising their cognition, which is often nuanced and context-dependent. This suggests that what students respond in written tests should be considered only as what they were triggered to write on that day in response to those particular questions, and may not fully reflect their knowledge and understanding of science topics. Other slightly different questions may well have cued the elicitation of different knowledge. Now Umar had recalled that chlorine comprises of covalent molecules, I asked him about the nature of the bond:

So what would that be, covalently bonded?

They share the electrons.

So how many electrons would they have then?

They’ll have

[Pause, c.7s – n.b., quite a long pause]

like the one on it, the one of the chlorines shares electrons with the other chlorine to fill in its shell on the other one, and the same does it with the other.

In thinking about covalent bonding, Umar (in common with many students) drew upon the full shells explanatory principle that considered bonding to be driven by the needs of atoms to 'fill' their outer electron shells. (The outer shell of chlorine would only actually be 'full' with 18 electrons, but that complication is seldom recognised, as octets and full shells are usually considered synonymous by students).

So how many electrons does each chlorine have to start with?

In the outer shell, seven.

And how many have they got after this?

They’ve got seven, but they share one.

[Pause, c.1s]

Maybe.

So that’s a covalent bond, is it?

Yeah.

So how many electrons are involved in a covalent bond?

[Pause, c.3s]

Erm,

[Pause, c.3s]

Two.

Two electrons.

So where do those two electrons come from?

They like, one that fills up the gap, fills up the – last electron needed in one of the chlorine shells, and the other chlorine shell fills it up in the other one.

So where do they come from?

Each chlorine. Outer shell.

One from each chlorine?

Yeah.

Okay, and that’d be a covalent bond?

Yeah.

Here, again, Umar is using the full shells explanatory principle as the basis for explaining the bond in terms of electrons 'filling up the gaps' in the electron shells, rather than considering how electrical interactions can hold the structure together. Umar's suggestion that the sharing of electrons "fills up the – last electron needed in one of the chlorine shells" demonstrates the anthropomorphic language (e.g., what an atom wants or needs) commonly used when learners have acquired aspects of the common octet rule framework that is developed from the full shells explanatory principle and used by many learners to explain bonding reactions, chemical reactions, patterns in ionisation energy, and chemical stability.

Sodium and chlorine don't actually overlap or anything

Keith S. Taber

Annie was a participant in the Understanding Chemical Bonding project. She was interviewed near the start of her college 'A level' course (equivalent to Y12 of the English school system). Annie was shown, and asked about, a sequence of images representing atoms, molecules and other sub-microscopic structures of the kinds commonly used in chemistry teaching. She was shown a representation of part of a lattice in sodium chloride.

Focal figure (Fig. 5) presented to Annie

Any idea what that's meant to be?

(pause, c.6s)

Just sodium and chlorine atoms

That's sodium and chlorine atoms, erm would you say that there was any kind of bonding there?

No.

Although the image included the standard '+' and '-' symbols to signify that ions were shown, Annie referred to "atoms". It transpired that Annie had an idiosyncratic understanding of what was meant by charge. (Read: Na+ has an extra electron in its outer shell and Cl- is minus an electron and K-plus represents a potassium atom that has an extra electron.)

Annie had already identified chemical bonding in representations of molecules of hydrogen , tetrachloromethane , and oxygen, so she was asked why she though there was no bonding in this example:

No bonding. Why do you say that? What is the difference between that and the ones we've seen before?

Well the other ones electrons were shown, and these no electrons are shown and they don't actually overlap or anything they just go in rows.

They go in rows. Okay. … but unlike (the images) we've seen previously they've had bonds in,

Yeah.

chemical bonds, whereas this, we don't have chemical bonds?

No.

So Annie did not interpret the representation of NaCl as portraying bonding. However, on further probing she did recognise that the structure could get held together by forces.

When Annie was asked if what was shown in the figure would would fall apart or hold together, Annie suggested that If you heated it, or reacted it in some way, it would hold together, and it would probably get held together by just forces. However, she did not consider that (i.e., even after reacting) amounted to chemical bonding. (Read: Sodium has one extra electron in its outer shell, and chlorine is minus an electron, so by force pulls they would hold together.)

The canonical interpretation of the figure is that it is a slice through a three dimensions structure of ions, where the attractive forces between cations pull the ions into a bound structure (to the point where attraction and repulsions are in equilibrium), and that this kind of binding is called ionic bonding.

Annie did not see ions, but atoms. She thought there was no bonding because no overlap was shown. In chemistry a wide range of different types of representation are used to show structures at the submicroscopic level – bonds may sometimes be shown by lines or sometimes by overlap or (in the case of ionic structures) neither. This is a potential source of confusion for learners who may not appreciate why different conventions may be used to represent different, or even the same, structures.

Ionic bonding – where the electron's transferred to complete the outer shell

Keith S. Taber

Amy was a participant in the Understanding Science Project. The first time I talked to Amy, near the start of her GCSE 'triple science' course in Y10 she told me that "in normal chemistry (i.e., the chemistry part of 'double science', as opposed to the optional additional chemistry lesson as part of 'triple science' that Amy also attended) we're doing about ionic bonding" which was "atoms which have either lost or gained electrons so they are either positively or negatively charged" and

"how the outer electron's transferred…to complete the outer shell of the erm chlorine, thing, ion…and the sodium atom loses erm, one electron is it, yeah one electron, erm, which the chlorine atom gains, and that yeah that completes its outer shell and makes the sodium positively charged and the chlorine negatively charged".

Amy told me that "in ionic bonding it's the electrons that are transferred, I think."

So Amy had acquired a common alternative conception, i.e. that ionic bonding involved electron transfer, and that this occurs to atoms to complete their electron shells.

Ionic bonding refers to the forces between ions that hold the structure of an ionic substance together, rather than a mechanism by which such ions might hypothetically be formed – yet often learners come away form learning about ionic bonding identifying it with a process of electron transfer between atoms instead of interactions between ions which can be used to explain the properties of ionic substances.

Moreover, the hypothetical electron transfer is a fiction. In the case of NaCl such an electron transfer between isolated Na and Cl atoms would be energetically unfavourable, even if reactants containing discrete atoms were available (which is unrealistic).

Whether students are taught that ionic bonding is electron transfer is a moot point, but often introductory teaching of the topic focuses not on the nature of the bonding, but on presenting a (flawed) teaching model of how the ions in the ionic structure could form by electron transfer between atoms. As this mechanism is non-viable, and so not an authentic scientific account, it may seem odd that teachers commonly offer it.

One explanation may simply be custom or tradition has made this an insidious alternative conception. Science teachers and textbooks have 'always' offered the image of electron transfer as representing ionic bonding. So, this is what new teachers had themselves been taught at school, is what they often see in textbooks, and so what they learn to teach.

Another possible explanation is in terms of what what is known as the atomic ontology. This is the idea that the starting pint for thinking about chemistry at the submicroscopic level is atoms. Atoms do not need to be explained (as if in nature matter always starts as atoms – which is not the case) and other entities such as ions and molecules do need to be explained in terms of atoms. So, the atomic ontology is a kind of misleading alternative conceptual framework for thinking about chemistry at the submicroscopic level.

A sodium atom wants to donate its electron to another atom

Keith S. Taber

Lovesh was a participant in the Understanding Chemical Bonding Project, studying 'A level' chemistry in a further education college. He was interviewed in his second year of the two year A level course, and was presented with focal figure 1 (below). He recognised figure 1 as showing a "sodium, atom", and was asked about its stability:

Is that a stable species, do you think?

Erm (pause, c.3s) No, because it hasn't got a, a full outer – electron shell, outer electron shell hasn't got eight electrons in.

Lovesh shared the common notion that an atom without a full outer shell / octet of electrons would be unstable compared with the corresponding ion with a full outer shell / octet of electrons. When comparing isolated atoms with the corresponding ions this is seldom the case, yet this is a common alternative conception about chemical stability. A sodium ion can be considered stable in an ionic lattice, or when hydrated in solution, but does not spontaneously ionise as the outer shell electron is attracted to the atom's positive core. Ionisation only occurs when sufficient work is done to overcome this attraction.

Lovesh was demonstrating the common full shells explanatory principle alternative conception which is central to the common octet rule framework – an alternative conceptual framework reflecting very common 'misconceptions' found among learners studying chemistry.

Lovesh was asked what would happen to the atom that he considered unstable:

So if it's not stable, what would tend to happen to that, do you think?

It will wanna donate the electron to another atom.

Right, when you say 'it wants to donate' it?

Erm. (pause, c.3s) Well because that outer electron is less attracted to the nucleus, erm it is, it can easily be transferred, attracted by another atom.

Lovesh's first response here used the term 'wanna' (want to) which if take literally suggests the atom has desires and preferences. This is an example of anthropomorphism, imbuing objects with human-like traits. Using anthropomorphic explanations is a common feature of the octet rule framework which often leads to students talking as if atoms deliberately act to get full outer electron shells.

It has been suggested that such anthropomorphism may be either 'strong'- where the learner is offering an explanation they find convincing – or 'weak' if they are using language metaphorically, just as a figure of speech.

In this case, when Lovesh's use of the notion of 'wants' was queried he was able to shift to a different language register in terms of the action of physics forces – the electron being attracted elsewhere. Lovesh had clearly acquired an appropriate way of thinking about the interactions between atoms, but his spontaneous explanation was couched in anthropomorphic terms. Although in this case the anthropomorphism was of a weak form, the habitual use of this kind of language may come to stand in place of offering a scientifically acceptable account.