Tag: chemical bonding

A tangible user interface for teaching fairy tales about chemical bonding

Keith S. Taber

Image by S. Hermann & F. Richter from Pixabay
Once upon a time there was a nometal atom that was an electron short of a full outer shell. "I wish I had an octet" she said, "if only I knew a nice metal atom that might donate their extra electron to me"… Image by S. Hermann & F. Richter from Pixabay

 

Today I received one of those internet notifications intended to alert you to work that you might want to read:

"You wrote the paper A common core to chemical conceptions: learners' conceptions of chemical…. A related paper is available on Academia.

Tangible interaction approach for learning chemical bonding"

an invitation to read
An invitation to read

I was intrigued. Learning (and teaching) about chemical bonding concepts has been a long-standing interest of mine, and I have written quite a lot on the topic, so I clicked-through and downloaded the paper.

The abstract began

"In this paper we present ChemicAble, a Tangible User Interface (TUI) for teaching ionic bonding to students of grade 8 to 10. ChemicAble acts as an exercise tool for students to understand better the concepts of ionic bonding by letting them explore and learn…."

Ionic bonding – an often mislearnt topic

This led to mixed feelings.

Anything that can support learners in making sense of the abstract, indeed intangible, nature of chemical bonding offered considerable potential to help learners and support teachers. Making the abstract more concrete is often a useful starting point in learning about theoretical concepts. So, this seemed a very well-motivated project that could really be useful.

It is sometimes argued that educational research is something of an irrelevance as it seldom impacts on classroom practice. In my (if, perhaps, biased) experienced, this is not so – but it is unrealistic to expect research to bring about widespread changes in educational practice quickly, and arguments that most teachers do not read research journals and so do not know who  initiated particular proposals has always seemed to me to be missing the point. We are not looking for teachers to pass tests on the content of research literature, and it is quite natural that the influence of research is usually indirect through, for example, informing teacher education and development programmes, or through revisions of curriculum, recommended teaching schemes, or formal standards.

This study by Agrawal and colleagues was not a theoretical treatise but a report of the implementation of a tool to support teaching and learning – the kind of thing that could directly impact teaching. So this was all promising.

However,I  also knew only too well that ionic bonding was a tricky topic. When I started research into learners' developing understanding of chemical bonding (three decades ago, now) I read several studies suggesting there were common alternative conceptions, that is misunderstandings, of ionic bonding found among students (e.g., Butts & Smith,  1987).

My own research suggested these were not just isolated notions, but often reflected a coherent alternative conceptual framework for ionic bonding that I labelled the 'molecular' framework (Taber, 1994, 1997). Research I have seen from other contexts since, leads me to believe this is an international phenomenon, and not limited to a specific curriculum context (Taber, 2013).

(Read about 'the Understanding Chemical Bonding project')

Ionic bonding – an often mistaught topic?

Indeed, I feel confident in suggesting:

  • secondary level students very commonly develop an alternative understanding of ionic bonding inconsistent with the scientific account…
  • …which they find difficult to move beyond should they continue to college level chemistry…
  • … and which they are convinced is what they were taught

Moreover, I strongly suspect that in quite a few cases, the alternative, incorrect model, is being taught. It is certainly presented, or at least implied, in a good many textbooks, and on a wide range of websites claiming to teach chemistry. I also suspect that in at least some cases,  teachers are teaching this, themselves thinking it is an acceptable approximation to the scientific account.

(Read about 'The molecular framework for ionic bonding')

A curriculum model of ionic bonding

So, I scanned the paper to see what account of the science was used as the basis for planning this teaching tool. I found this parenthetical account:

"{As stated in the NCERT book on Science for class X, chapter 3, 4, the electrons present in the outermost shell of an atom are known as the valence electrons. The outermost shell of an atom can accommodate a maximum of 8 electrons. Atoms of elements, having a completely filled outermost shell show little chemical activity. Of these inert elements, the helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell.

The combining capacity of the atoms of other elements is explained as an attempt to attain a fully-filled outermost shell (8 electrons forming an octet). The number of electrons gained, lost or shared so as to make the octet of electrons in the outermost shell, gives us directly the combining capacity of the element called the valency. An ion is a charged particle and can be negatively or positively charged. A negatively charged ion is called an 'anion' and the positively charged ion, a 'cation'. Metals generally form cations and non-metals generally form anions. Atoms have tendency to complete their octet by this give and take of electron forming compounds. Compounds that are formed by electron transfer from metals to non-metals are called ionic compounds.}"

Agrawal et al., 2013 (no page numbers)

There are quite a few ideas here, and quite a lot of his account is perfectly canonical, at least at the level of description suitable for secondary school, introductory, chemistry. However, sprinkled in are some misleading statements.

So,

Curriculum statement Commentary
"…the electrons present in the outermost shell of an atom are known as the valence electrons."

 Fine
"The outermost shell of an atom can accommodate a maximum of 8 electrons."

This is only correct for period 2.

It is false false for period 1 (2 electrons), period 3 (18 electrons), period 4 (32 electrons), etcetera.
"Atoms of elements, having a completely filled outermost shell show little chemical activity. Of these inert elements, the helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell."

Fine – apart from the reference to  "completely filled outermost shell"

Of the noble gases, only helium and neon have full outer shells.

'Atoms' of the heavier noble gases with full outer shells would not atoms, but ions, and these would be extremely unstable – i.e., they could not exist except hypothetically under extreme conditions of very intense electrical fields.
"The combining capacity of the atoms of other elements is explained as an attempt to attain a fully-filled outermost shell (8 electrons forming an octet). The number of electrons gained, lost or shared so as to make the octet of electrons in the outermost shell, gives us directly the combining capacity of the element called the valency."

Hm –  generally the valency can be identified with the difference between an atom's electronic configuration and the 'nearest' noble gas electronic configuration – which would be an octet of valence shell electrons, except in period one.

However,  the equivalence suggested here "a fully-filled outermost shell (8 electrons forming an octet)" is only true for period 2. An octet does not suffice for a full outer shell in period 3 (full at 18  electrons), or in period 4 (full at 32 electrons), etcetera.

And, in the statement, valency is described as being related to the intentions of atoms: "is explained as an attempt to attain…" (and "…electrons gained, lost or shared so as to…") which encourages student misconceptions. [Read about 'Learners' anthropomorphic thinking'.]
"An ion is a charged particle and can be negatively or positively charged. A negatively charged ion is called an 'anion' and the positively charged ion, a 'cation'. Metals generally form cations and non-metals generally form anions." Fine.
"Atoms have tendency to complete their octet by this give and take of electron forming compounds."

This is a common notion, but actually suspect. Some elements have an electron affinity such that the atoms would tend to pick up an electron spontaneously.

However, for an element with a valency of -2, such as oxygen, once it has become a singly charged anion (O–), it will not attract a second electron, so apart from the halogens, this is misleading. The negatively charged O– ion will indeed spontaneously repel/be repelled by a (negatively charged) electron.

Metallic elements have ionisation enthalpies showing that energy has to be applied to strip electrons from them – they certainly do not have a "tendency to complete their octet by this giv[ing]" of electrons.
"Compounds that are formed by electron transfer from metals to non-metals are called ionic compounds."

This is not usually how ionic compounds are formed. Although it is possible in the lab. to use binary synthesis (e.g., burning sodium in chlorine – not for the faint-hearted), that is not how ionic compounds are prepared in industry, or how the NaCl in table salt formed naturally.

(And even when burning sodium in chlorine, neither of the reactants are atomic, so even here there is no simple transfer of electrons between atoms.)

So this account is a mixture of the generally correct; the potentially misleading; and the downright wrong.

Agrawal and colleagues describe an ingenuous apparatus they had put together so that students can physically manipulate tokens to see ionic bond formation represented. This looks like something that younger secondary children would really enjoy.

They also report a small-scale informal evaluation of a classroom test of the apparatus with an unspecified number of students, reporting very positive responses. The children generally found the apparatus easy to use, the information it represented easy to understand, and they thought it helped them learn about chemical [ionic] compound formation.  So this seems very successful.

However, what did it help them learn?

The teaching model

"For example, when a token representing [a] sodium atom is placed on the table top, its valence shell (outermost shell) with 1 revolving valence electron is displayed around the token. When the student places a chlorine atom on the table, its valence shell along with 7 revolving valence electrons is displayed. The electron from the sodium atom gets transferred to the chlorine atom. +1 charge appears on the sodium atom due to loss of electron and -1 charge appears on the chlorine atom due to gain of electron. Both form a stable compound. The top bar on the user interface turns green to show success and displays the name of the stable compound so formed (sodium chloride, in this case). The valence shell of the atoms also turns green to show a stable compound."

Agrawal et al., 2013 (no page numbers)

Which sounds impressive, except NaCl is not formed by electron transfer, and with the ChemicAble the resulting structure is a single Na+-Cl– ion pair, which does not represent the structure of the NaCl compound, and indeed would not be a stable structure.

Does it matter if children are taught scientific fairy tales?

The innovation likely motivated learners. And the authors seem to be basing their 'ChemicAble' on the curriculum models set out in the model science books produced by the Indian National Council of Educational Research and Training. So, the authors have produced something that helps children learn the science curriculum in that context,and so presumably what students will subsequently be examined on. Given that, it seems churlish to point out that what is being taught is scientifically wrong.

So, I find it hard to be critical of the authors, but I do wonder why governments want children to learn scientific fairy tales that are nonsense. The electron transfer model of ionic bonding seems to be popular with teachers, and received well by learners, so if the aim of education is to find material to teach that we can then test children on (so they can be graded, rated, sequences, selected), what is the problem? After all, I am a strong advocate for the idea that what we teach in school science is usually, necessarily, a simplification of the science – and indeed is basically a set of models – and not some absolute account of the universe.

Here the children, the teacher and the researchers have all put a lot of effort into helping learners acquire a scientifically incorrect account of ionic bonding. We think children should learn about the world at the molecular, naometre scale as this is such an important part of chemistry as a science. Yet, to my mind, if we are going to ask children to put time and effort into learning abstract models of the structure of nature at submicroscopic levels, even though we know this is challenging for them, then, although we need to work with simplified models, these should at least be intellectually honest models, and not accounts that we know are completely inauthentic and do not reflect the science. This is why I have been so critical of the incoherence and errors in the chemistry in the English National Curriculum (Taber, 2020).

Otherwise, education is reduced to a game for its own sake, and we may as well ask students to learn random Latin texts, or the plots of Grimms' Fairy Tales, or even the chemical procedures obscured by disguised reagents and allegorical language in alchemical texts, and then test them on how much they retain.

Actually, no, this learning of false models is worse than that, because learning these incorrect accounts confuses students and impedes their learning of the canonical scientific models if they later go on to study the subject further. So, if it is important that children learn something about ionic bonding, let's teaching something that is scientifically authentic and stop offering fairly tales about atoms wanting to fill their shells.

Sources cited:
 
 

 

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.

Single bonds are different to covalent bonds

Single bonds are different to covalent bonds or ionic bonds

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 commonl y used in chemistry teaching. She was shown a representation of the resonance between three canonical forms of BF3, sometimes used as away of reflection polar bonding. She had just seen another image representing resonance in the ethanoate ion, and had suggested that it contained a double bond. She had earlier in the interview referred to covalent bonding and ionic bonding, and after introducing the ideas of double bond, suggested that a double bond is different to a covalent bond.

Focal figure (14) presented to Annie

What about diagram 14?…

Oh.

(pause, c.13s)

Seems to be different arrangements. Of the three, or two elements.

Uh hm.

(pause, c.3s)

Which are joined by single bonds.

What, where, what single, what sorry are joined by single bonds?

All the F to the B to the F. Are single bonds they are not double like before. [i.e., a figure discussed earlier in the interview]

So are they covalent bonds? Or ionic bonds, or? Or are single bonds something different again?

Single bonds are different.

This reflected her earlier comment to the effect that a double bond is different to a covalent bond, suggesting that she did not appreciate how covalent bonds are considered to be singular or multiple.

However, as I checked what she was telling me, Annie's account seemed to shift.

They're different to double bonds?

Yeah.

And are they different to covalent bonds?

No 'cause you probably get covalent bonds which are single bonds.

So single bonds, just moments before said to different to covalent bonds, were now 'probably' capable of being covalent. As she continued to answer questions, Annie decided these were 'probably' just alternative terms.

So covalent bonds and single bonds, is that another word for the same thing?

Yeah, probably. But they can probably occur in different, things like in organic you talk about single bonds more than you talk about covalent, and then like in inorganic you talk about covalent bond, more than you talk about single bonding or double bonding.

So you think that maybe inorganic things, like sort of, >> copper iodide or something like that, that would tend to be more concerned with covalent bonds?

< Yeah. < Yeah.

But if you were doing organic things like, I don't know, erm, ethane, >> that's more likely to have single bonds in.

< Yeah. < Yeah.

So single bonds are more likely to occur in carbon compounds.

Yeah.

And covalent bonds are more likely to occur in some other type of compound?

Yeah. Sort of you've got different terminology, like you could probably use single bonds to refer to something in inorganic, but when you are talking about the structures and that, it's easier to talk about single bonds and double bonds, rather than saying that's got a covalent bond or that's got an ionic bond.

Annie's explanation did not seem to be a fully thought-out position. It was not consistent with the way she had earlier reported there being five covalent bonds and one double bond in an ethanoate ion.

It seems likely that in the context of the research interview, where being asked directly about these points, Annie was forced to make explicit the reasons she tended to label particular bonds in specific ways. The interview questions may have acted like Socratic questioning, a kind of scaffolding, leading to new insights. Only in this context did she realise that the single and double bonds her organic chemistry lecturer talked about might actually be referring to the same entities as the covalent bonds her inorganic chemistry lecturer talked about.

It would probably not have occurred to Annie's lecturers (of which, I was one) that she would not realise that single and double bonds were covalent bonds. It may well have been that if she had been taught by the same lecturer in both areas, the tendency to refer to single and multiple bonds in organic compounds (where most bonds were primarily covalent) and to focus on the covalent-ionic dissension in inorganic compounds (where degree of polarity in bonds was a main theme of teaching) would still have lead to the same confusion. Later in the interview, Annie commented that:

if I use ionic or covalent I'm talking about, sort of like a general, bond, but if I use double or single bonds, that's mainly organic, because sort of it represents, sort of the sharing, 'cause like you draw all the molecules out more.

This might be considered an example of fragmentation learning impediment, where a student does not make a link that the teacher is likely to assume is obvious.

Covalent bonding is sharing electrons

It's covalent bonding where the electrons are shared to create a full outer shell

Keith S. Taber

Brian was a participant in the Understanding Chemical Bonding project. He was interviewed during the first year of his college 'A level' course (equivalent to Y12 of the English school system). Brian 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. He was shown a simple representation of a covalent molecule:

Focal figure ('2') presented to Brian

Any idea what that's meant to be, number 2?

Hydrogen molecule.

Why, how do you recognise that as being a hydrogen molecule?

Because there's two atoms with one electron in each shell.

Uh hm. Er, what, what's going on here, in this region here, where these lines seem to meet?

Bonding.

That's bonding. So there's some sort of bonding there is there?

Yeah.

Can you tell me anything about that bonding?

It's covalent bonding.

So, so what's covalent bonding, then?

The electrons are shared to create a full outer shell.

Okay, so that's an example of covalent bonding, so can you tell me how many bonds there are there?

One.

There's one covalent bond?

Yeah.

Right, what exactly is a covalent bond?

It's where electrons are shared, almost, roughly equally, between the two atoms.

So that's what we'd call a covalent bond?

Yeah.

So according to Brian, covalent bonding is where "the electrons are shared to create a full outer shell". The idea that a covalent bond is the sharing of electrons to allow atoms to obtain full electron shells is a very common way of discussing covalent bonding, drawing upon the full shells explanatory principle, where a 'need' for completing electron shells is seen as the impetus for bonding, reactions, ion formation etc. This principle is the basis of a common alternative conceptual framework, the octet rule framework.

For some students, such ideas are the extent of their ways of discussing bonding phenomena. However, despite Brian defining the covalent bond in this way, continued questioning revealed that he was able to think about the bond in terms of physical interactions

Okay. And why do they, why do these two atoms stay stuck together like that? Why don't they just pull apart?

Because of the bond.

So how does the bond do that?

(Pause, c.13s)

Is it by electrostatic forces?

Is it – so how do you think that works then?

I'm not sure.

The long pause suggests that Brian did not have a ready formed response for such a question. It seems here that 'electrostatic forces' is little more than a guess, if perhaps an informed guess because charges and forces had features in chemistry. A pause of about 13 seconds is quite a lacuna in a conversation. In a classroom context teachers are advised to give students thinking time rather than expecting (or accepting) immediate responses. Yet, in many classrooms, 13 seconds of 'dead air' (to borrow a phrase from broadcasting) from the teacher night be taken as an invitation to retune attention to another station.

Even in an interview situation the interviewer's instinct may be to move on to a another question, but in situations where a researcher is confident that waiting is not stressful to the participant, it is sometimes productive to give thinking time.

Another issue relating to interviewing is the use of 'leading questions'. Teachers as interviewers sometimes slip between researcher and teacher roles, and may be tempted to teach rather than explore thinking.

Yet, the very act of interviewing is an intervention in the learners' thinking, in that whatever an interviewer tells us is in the context of the conversation set up by the interviewer, and the participant may have ideas they would not have done without that particular context. In any case, learning is not generally a once off event, as school learning relies on physiological process long after the initial teaching event to consolidate learning, and this is supported by 'revision'. Each time a memory is reactivated it is strengthened (and potentially changed).

So the research interview is a learning experience no matter how careful the researcher is. Therefore the idea of leading questions is much more nuanced that a binary distinction between those questions which are leading and those that are not. So rather than completely avoiding leading questions, the researcher should (a) use open-ended questions initially to best understand the ideas the learner most easily beings to mind; (b) be aware of the degree of 'scaffolding' that Socratic questioning can contribute to the construction of a learners' answer. [Read about the idea of scaffolding learning here.] The interview continued:

Can you see anything there that would give rise to electrostatic forces?

The electrons.

Right so the electrons, they're charged are they?

Yeah. Negatively.

Negatively charged – anything else?

(Pause, c.8s)

The protons in the nucleus are positively charged.

Uh hm. And so would that give rise to any electronic interactions?

Yeah.

So where would there be, sort of any kind of, any kind of force involved here is there?

By the bond.

So where would there be force, can you show me where there would be force?

By the, in the bond, down here.

So the force is localised in there, is it?

The erm, protons would be repelling each other, they'd be attracted by the electrons, so they're keep them at a set distance.

It seemed that Brian could discuss the bond as due to electrical interactions, although his initial ('instinctive') response was to explain the bond in terms of electrons shared to fill electron shells. Although the researcher channelled Brian to think about the potential source of any electrical interactions, this was only after Brian had himself conjectured the role of 'electrostatic forces.'

Often students learn to 'explain' bonds as electron sharing in school science (although arguably this is a rather limited form of explanation), and this becomes a habitual way of talking and thinking by the time they progress to college level study.

They're both attracting each other but this one's got a larger force

Iodine's got a larger force that lithium, so it will pull towards the lithium more 

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.

When she was shown an image representing the electron cloud around an iodide ion polarised by an adjacent lithium ion Annie interpreted this as the iodine exerting a greater force on the lithium than vice versa.

Focal figure presented to Annie

What about this, any idea about this?

It's the same sort of thing again – the lithium combines with the iodine – to make a stable outer shell between the two, by sharing electrons, but the lithium has a smaller charge, or smaller pull than the iodine, so the actual shape of it goes in towards. It sort of goes inwards because its attracting the lithium, whereas if the lithium was attracting it, it would be like a reverse picture.

So, so the iodine's attracting what, sorry?

The lithium.

The iodine's attracting the lithium, and the lithium is not attracting the iodine?

Yeah, they're both attracting each other but because this one's got a larger force, then it will pull towards the lithium more.

The iodine's got a larger force,

Yeah.

so it will pull towards the lithium more?

Yeah.

Any image used to represented chemical bonding is necessarily a kind of model, and a partial representation – and there are a range of types of representations students meet. It is perhaps not surprising if students cannot always 'guess what the teacher (or textbook author or researcher) is thinking, and what they intend by a particular type of image.

Annie here demonstrates the common notion that chemical bonding can be based upon 'sharing' electrons (i.e., covalent bonding). At this point in her course Annie would not be expected to appreciate polar bonds or the polarisation of ions, but her prior learning that covalent bonding could be understood as 'sharing' of electrons could potentially act as an impediment to learning that the ionic-covalent bonding distinction should be seen as a spectrum, a continuous dimension, not a dichotomy.

The way forces are understood in physics is that they are interactions between two bodies, and that the same magnitude of force acts of both bodies (i.e., Newton's third law). However, students commonly consider that a 'larger' body (e.g., more massive, more highly charged) exerts a large force on the smaller body. Students do not clearly distinguish the force from its effect, and so this alternative conception seems to draw upon intuitions based on actual experience of the world (i.e., a grounded learning impediment) where larger sources (larger fires, bigger loudspeakers, larger lamps) often seem to have larger effects.

[Read about Newton's third law, and student learning difficulties]

In a molecule, the electron actually slots into spaces

Keith S. Taber

Mohammed was a participant in the Understanding Science Project. When interviewed in the first term of his upper secondary (GCSE) science course (in Y10), he told me he had been learning about ionic bonding in one of his science classes. Mohammed had quite a clear idea about ionic bonding, which he described in terms of the interactions of two atoms where "they both want to get full outer shells", leading to salt which was "like two atoms joined together":

The "two atoms joined together" sounds much like a molecule (and it is very common for students to identify molecule like ion-pairs even in representations of extensive ionic lattices), so I asked Mohammed about this:

Can I see these atoms?

No. They're really small. Because the wavelength of visible light is actually too like large to see the atoms, they just pass over them.

Okay, so I can't see them. But I can imagine them, can I?

Yeah.

So if I could imagine a sodium atom and chlorine atom, and then they form salt, what would it look like afterwards? How could I imagine it afterwards.

Oh it's like two atoms joined together.

That sounds like a molecule to me?

It's not actually, like, joined.

No?

Because I know that whenever things of opposite charge, I know two rods, when they come together, they don't actually touch, so they don't exactly touch, but they are very close, two atoms close to each other

So a molecule would be different to that in some way, would it?

Yeah, a molecule's actually bonded

So how that different?

I think in a molecule, the electron actually slots into spaces.

I see, and it doesn't do that in this case?

No.

So Mohammed thinks that the interaction between the ions will be due to their electrical charges, but, for him, this may not count as a bond, as the forces just hold the ions ("atoms") close together, and do not actually join them. Mohammed's idea of the atoms not actually touching, "they don't actually touch, so they don't exactly touch", is transferring a notion from the familiar world of macroscopic phenomena (where things touch, or they do not touch) to the submicroscopic world of quanticles that do not have definitive size/volume, and do not actually have distinct surfaces, so touching is a matter of degree. There is no more (or less) 'touching' in a covalent bond than in ionic bonding. So according to Mohammed the ions do not form a molecule, as in a molecule there would some kind of more direct joining – he suggests something like an interlocking with electrons from one atom slotting into spaces on another.

Interestingly, Mohammed bases his notion that the ions would not touch on a general principle that he considers to apply whenever considering things of opposite charge – which he justifies on his knowledge that "two [charged] rods, when they come together, they don't actually touch". He may be misremembering something here – or he may have seen a demonstration of suspended charged rods of the same material (so either both negatively or both positively changed) that when one is moved closer to the other the rods repel. Whatever the source, Mohammed seems to feel he has a valid general principle that he can apply here that act as a grounded learning impediment channelling his thinking about the case under discussion along 'the wrong lines'.

Mohammed's notion of the ionic bonding as being just due to forces rather than being a proper bond is very similar to a common alternative conceptions of ionic bonding which sees ions in a lattice only having a limited number of ionic bonds depending upon valency (the valency conjecture) but bonded with other coordination counter-ions by 'just forces' (the just forces conjecture) – although here Mohammed suspected that all ionic bonding fell short of being proper chemical bonds.

This is a very mechanical model of the covalent bond, whereas the scientific model presents bonding as more of a process than a material mechanical link. However teaching models often present bonding this way, and sometimes molecules are modelled in terms of jigsaws with atoms or radicals as pieces to be slotted together. Although such models are only meant to provide a simple analogy for the bonding they may act as learning impediments if learners take them too 'literally' as realistic representations and transfer inappropriate associations from the model to their understanding of the system being modelled.

Mohammed also uses similar language when asked about salt dissolving in water, as the charge of the water forces the sodium and chlorine ions to slot into certain places within the water molecules *.

Salt is like two atoms joined together

Keith S. Taber

Mohammed was a participant in the Understanding Science Project. When interviewed in the first term of his upper secondary (GCSE) science course (in Y10), he told me he had been learning about ionic bonding in one of his science classes. Mohammed had quite a clear idea about ionic bonding, which he described in terms of the interactions of two atoms where "they both want to get full outer shells"*.

Some learners have been found to see the electron transfer process described by Mohammed (which is purely a way of conceptualising ion formation, and has little connection with what actually happens when ionic substances form) as being the bond, and sufficient to hold species together. However, Mohammed did recognise the role of electrical forces in holding the species together:

And did you say that if you take a sodium atom and a chlorine atom, you get salt?

Yeah… Sodium chloride. And the way they bond together, is because now one, the sodium has lost an electron. And they start off neutral because the protons and electrons balance each other out, because the same number of them, but when you lose one you get plus one in the sodium, and like when you have a chlorine, you add an electron so you get minus one. In the end the whole compound is neutral, but because erm, like they, they're differently charged they attract together, and they bond together. I think.

So if I could imagine a sodium atom and chlorine atom, and then they form salt, what would it look like afterwards? How could I imagine it afterwards.

Oh it's like two atoms joined together.

So here Mohammed is clear (if somewhat tentative) that a bond has been formed. Yet his focus is on the iteration of two atoms, forming ions, whereas ionic bonding needs to be understood as the various interactions at work in a lattice. The process of ion formation described by Mohammed: 

  • would not actually be energetically viable for two atoms (as the electron affinity of chlorine is 364 kJ mol-1, whereas the first ionisation energy of sodium is 494 kJ mol-1, so the electron capture of a chlorine atom would not release enough energy to remove the electron from the sodium atom);
  • neither sodium nor chlorine atoms are stable under normal chemical conditions: neither sodium nor chlorine used in a binary synthesis would be in the form of discrete atoms; and sodium chloride is more likely to be formed by neutralisation using substances where the sodium and chloride ions are already present, e.g. sodium hydroxide and hydrochloric acid.)

Mohammed's explanation conflates two levels – the macroscopic level of bench phenomena (such as the substance sodium chloride – common salt) and the level of models at the submicroscopic scale of molecules, ions and atoms. Even if one atom of sodium could interact ('react' is better kept as a term for what occurs at the level of substances) with one atom of chlorine in the manner Mohammed envisages, to give the "two atoms joined together", that entity could not meaningfully be identified as salt as many of the properties of salts emerge from the ionic lattice of myriad ions.

[If sodium-chloride ion pairs could be formed then we might consider these as the component quanticles of a form of sodium chloride, but this would need to be considered a different allotrope to table salt, just as ozone molecules are not the basis of the oxygen in air that supports respiration.]

The "two atoms joined together" sounds much like a molecule (and it is very common for students to identify molecule like ion-pairs even in representations of extensive ionic lattices), so I asked Mohammed about this.

Electrons would contain some of the element

Electrons from different elements would be different – perhaps because they would actually contain some of the element in the electron?

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). She was shown a representation of a tetrachlomethane molecule.

Understanding Chemical Bonding project – Focal figure 3

When Annie was asked about the diagram, she noted that (following a representational convention) the electrons were represented differently. Using different symbols like this is quite common, but is little more that a bookmaking tool – to help keep count of the number of electrons in the molecule in relation to those that would be present in discrete atoms.

…are there any bonds [shown] in that diagram do you think?

Yes.

How many?

Four.

Four bonds, so we've got four bonds there. Erm, are the bonds actually shown?

Yeah.

So how are they represented on the diagram?

By the circles that overlap, and they're showing it by the electrons, the outer-shell electrons in the chlorine have got black dots and the ones from carbon have got just circles.

Okay. So the carbon electrons and the chlorine electrons are signified in a different way

Yeah.

I followed up this point to check Annie understood that the convention did not imply that there was any inherent difference between the electrons.

So what would be the difference between a carbon electron and a chlorine electron?

(pause, c.5s)

The expected answer here was 'no difference', but the pause suggested Annie was not clear about this. So I set up an imaginary scenario, a kind of thought experiment:

If I gave you a bottle of electrons – which I can't do – how would you be able to tell chlorine electrons from carbon electrons – in what ways would they be different?

They would be different because, erm, I don't know if they would actually contain some of the element in the electron.

Do you think they might have little labels on some with "C"s and some with "Cl"s or

Yeah, I don't know if you got an electron, and you could sort of if you took one single one you could say, right that's chlorine and that one's carbon.

You are not sure, you are not sure if you could, or not?

No.

The idea that an electron might contain some of the element seems to miss the key idea that macroscopic phenomena (samples of element) are considerer to energy from extensive ensembles of submicroscopic particles ('quanticles').

Annie did not seem too sure here – perhaps her intuition was that a carbon electron would be different to a chlorine electron, but she could not suggest how. Electrons have no memories, and there is no way of knowing whether an electron has previously been part of a particular atom (or ion or molecule). A free electron is not meaningfully a chlorine electron or a carbon electron. However, students do not always appreciate this, and may consider that free electrons in some sense belong to an atoms they they derived form, and even that this may later have consequences (as with the 'history' conjecture in thinking about ionic bonding).

Annie went on to suggest that carbon electrons would be bigger than chlorine electrons.

A chemical bond would have to be made of atoms

Keith S. Taber

Amy was a participant in the Understanding Science Project. When I had talked to Amy when she was in Y10 she had referred to things being bonded: "where one thing is joined on to another thing, and it can be chemically bonded" and how "in a compound, where two or more elements are joined together, that's an example of chemical bonding".

The following year, in Y11, when she was studying fats she talked about "how they're made up and like with all the double bonds and single bonds" where a double bond was "where there are kind of like two bonds between erm carbon atoms instead of like one" and a bond was "how two atoms are joined together". Later in Y11, Amy told be that she did not know how to explain chemical bonding, but "in lessons like we've always been shown these kind of – things – where you kind of, you've got the atom, and then you've got the little, grey stick things which are meant to be the bonds, and you can just – fit them together."

Source: Image by WikimediaImages from Pixabay

As Amy had told me "everything is made up of atoms", I provocatively asked her if the chemical bond was made of atoms. Amy had "absolutely no idea" but she "suppose(d) it would have to be, wouldn't it".

Not only is this an alternative conception, but to a chemist, or science teacher, the idea that chemical bonds are themselves made up of atoms seems incongruous and offers a potential for infinite regress (are those atoms in the bonds, themselves bonded? If so, are those bonds also made of atoms?)

This alternative conception could be considered a kind of associative learning impediment – that is where a learner makes an unintended link and so applies an idea outside of its range of application. All material is considered to be made of atoms – or at least quanticles comprising one of more nuclei bound to electrons (i.e., ions, molecules). Even this is not an absolute: the material formed immediately after the big bang was not of this form, and nor is the matter in a neutron star, but the material we usually engage with is considered to be made of atom-like units (i.e., ions, molecules).

But to suggest that Amy has made an inappropriate association seems a little unfair. Had Amy thought "all matter was made of atoms" and then suggested that chemical bonding was made of atoms this would be inappropriate as chemical bonding is not material but a process – electrical interactions between quanticles. Yet it is hard to see how one can over-extend the range of 'everything', as in "everything is made up of atoms".

There is an inherent problem with the motto everything is made up of atoms. It is probably something that teachers commonly say, and think is entirely clear – that it is obvious what its scope is – but from the perspective of a student there is not the wealth of background knowledge to appreciate the implied limits on 'everything'.

Learners will readily pick up teaching mottos such as "everything is made of atoms" and take them quite literally: if everything is made of atoms then bonds must be made of atoms. So although she was wrong, I think Amy was just applying something she had learnt.

She'd never thought about whether ionic bonding is the same thing as chemical bonding

Keith S. Taber

Amy was a participant in the Understanding Science Project. When I talked to her near the start of her GCSE 'triple science' course in Y10 she told me that ionic bonding was "atoms which have either lost or gained electrons so they are either positively or negatively charged" and that chemical bonding was "like in a compound, where two or more elements are joined together", but she seemed unsure how the two concepts were related.

I followed up on Amy's use of the term 'compound' to explore how she understood the term:

How would you define a compound?

Erm …Something which has erm two or more elements chemically bonded.

… So you give me an example of that, compound?

Erm, sodium oxide.

Sodium oxide, okay, so there are two or more elements chemically bonded in sodium oxide are there?

Uh hm

And what would those two or more elements be?

Sodium and oxygen.

Okay. Erm, so when we say sodium oxide is chemically bonded, what we are saying there is?

[pause, c 2s]

Erm – a sodium atom has been bonded with a oxygen atom to form erm a new substance.

So Amy's example of a compound was sodium oxide, which would normally be considered essentially an ionic compound, that is a compound with ionic bonding. So this gave me an opportunity to test out whether Amy saw the bonding in sodium chloride and sodium oxide as similar.


Okay, so that was chemical bonding,

Mm.

and that occurs with compounds?

Yeah.

And what did you say about ionic bonding?

Erm, it's the outer electrons they are transferred from one element to another.

Now what does that occur in? You gave me one example, didn't you?

Uh huh

Sodium chloride?

Yeah
…

Erm. Would sodium chloride be er an element?

[pause, c.2s]

Sodium chloride, no.

No?

It would be a compound.

You think that would be a compound?

Yeah.

And a compound is two or more elements joined together by chemical bonding?

Yeah.

So Amy had told me that sodium chloride, which had ionic bonding, was (like sodium oxide) a compound, and she had already told me that a compound comprised of "two or more elements chemically bonded", so it should be follow that sodium chloride (which had ionic bonding) had chemical bonding.

Do you think sodium chloride has chemical bonding?

Er – I think so

And it also has ionic bonding, or is that the same thing?

Erm,

[pause, c.2s]

I dunno, I've never thought about it that way, erm,

[pause c.3s]

I'm not sure, erm

[pause, c.2s]

I dunno, it might be.

Clearly, whatever Amy had been taught (and interviewing students reveals they often only recall partial and distorted versions of what was presented in class) she had learnt

  • (1) that ionic bonding was transfer of electrons (an alternative conception) as in the example of sodium transferring an electron to chlorine; and that
  • (2) a compounds was where two or more elements chemically bonded together, and an example was sodium oxide where the elements sodium and oxygen were chemical bonded.

Yet these two pieces of learning seemed to have been acquired as isolated ideas without any attempt to link them. Initially Amy seemed to feel ionic bonding and chemical bonding were quite separate concepts.

When taken through an argument that led to her telling me that sodium chloride, that she thought had ionic bonding, was a compound, which therefore had chemical bonding, there should have been a logical imperative to see that ionic bonding was chemical bonding (actually, a kind of chemical bonding – as the logic did not imply that chemical bonding was necessarily ionic bonding). Despite the implied syllogism:

  • sodium chloride has ionic bonding
  • sodium chloride is a compound
  • compounds have elements chemically bonded together
  • therefore ionic bonding …

Amy was unsure what to deduce, presumably because she had seen the two concepts of ionic bonding and chemical bonding as discrete notions and had had given no thought to a possible relationship between them. However explicit teaching had been on this point, it is very likely that the teacher had expected students to appreciate that ionic bonding was a type of chemical bonding – but Amy had not integrated these ideas into a connected conceptual structure (i.e., there was a learning bug that could be called a fragmentation learning impediment).

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.

Sodium has one extra electron in its outer shell, and chlorine is minus an electron, so by force pulls they would hold together

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.

Focal figure (Fig. 5) presented to Annie

She was shown a representation of part of a lattice of ions in sodium chloride (see: Sodium and chlorine don't actually overlap or anything), but Annie identified the signified as atoms, not ions, because 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 was asked whether the structure made up of sodium and chlorine 'atoms' would hold together:

Do you think this thing would fall apart? Or would it hold together?

(pause, c.9s)

If you heated it, or reacted it in some way, it would hold together, and it would probably get held together by just forces.

By forces. Any idea what kind of forces would hold it together?

Probably just the attraction.

Uh hm?

The attraction from the plus to the minus because like chlorine's minus an electron and sodium is over an electron. So they could just like hold them together, but not actually combine.

Right, chlorine's, so sodium's, say that about the electrons again.

Sodium has like one extra electron, 'cause it has like an extra electron in its outer shell, and chlorine has seven electrons in its outer shell so its minus an electron so by sort of exchanging, the sodium combining with the chlorine just by force pulls they would hold together.

So Annie saw the plus (+) symbol to mean one electron over a full shell (2.8.1), and the minus (-) symbol to mean one electron short of an octet of electrons (2.8.7). For Annie these charges were not net electrical charges, but deviations from octet configurations. Yet, these 'deviation charges', for Annie, provided the basis for the attraction between the 'charged' atoms.

This was checked by asking Annie about the electron configurations.

So we looked at a sodium atom earlier, you recognised it as being a sodium atom, …

Can you tell me what the configuration is in terms of shells? How many in the first shell, how many in the second shell…

2.8.1

2.8.1?

Yeah.

So this here (indicating a cation on the figure), you are saying that this here is 2.8.1

Yes.

And this is 2.8.7 would it be?

Yeah, 2.8.7

And that is what holds them together the fact that this is one short,

yeah,

one over and one short.

One over, and that one's one short.

So the plus means one electron more than an outer, the full shell,

Yeah.

and the minus means one electron

Minus.

less than an outer shell,

Yeah.

and that's what holds them together.

Yeah.

Okay, so there is something holding them together,

right,

and it's to do with these pluses and these minuses,

Yes.

but what we don't have there is chemical bonding like we had before.

No.

Annie held an alternative conception of the nature of the charges associated with ions: that neutral atoms had 'charges' if they did not have full shells/octets of electrons. Whilst Annie's specific deviation charge conception would seem to be rather unusual, alternative conceptions relating to the significance of full shells / octets of electrons seems to be very common among chemistry students. Although Annie's thinking was idiosyncratic it reflected the common full shells explanatory principle that sees electronic configuration as a cause for chemical processes.

So Annie considered that these 'deviation' charges could actually give rise to forces between atoms (see also The force of lack of electrons pulls two hydrogen atoms together*).

Annie did not see ions, but atoms. But she thought that after a reaction, there would be attractions, 'force pulls', holding the product together, but this would not amount to chemical bonding.

Annie's notion of 'charges' on atoms (being extra or missing electrons in the outer shell), that led to her not recognising bonding in the NaCl, was an uncommon alternative conception notion. However, her notion that chemical bonding was something other than 'just forces', and that sometimes structures were held together by 'just forces' when there was no bonding, is a common alternative conception. Indeed it is part of a common 'molecular framework' for conceptualising ionic bonding, that is in turn a part of a common alternative conceptual framework for thinking about chemical bonding, stability and reactions: the octet framework.