Chapter 7 of Understanding Chemical Bonding: The development of A level students' understanding of the concept of chemical bonding
Stability and lability in cognitive structure: the case of Annie
§7.0: An overview of Annie’s case
Annie was one of the four colearners participating in the first stage of the interview study. She was interviewed on four occasions over a period of nearly 16 months: near the start of her second term of A level, near the end of her first year, and twice shortly before her A level examination (see appendix 1, §A1.1, for the schedule of interviews undertaken for the research). A case study was written around themes derived from the structure of the academic subject (see chapter 5, §5.3).
It was found that Annie's understanding of chemical bonding developed where her existing cognitive structure was labile enough for her to assimilate new ideas. However, the case study demonstrated that there were some aspects of Annie's cognitive structure that showed considerable stability, and where learning could only occur gradually.
It was also found that much of Annie's thinking about chemical bonding could be represented in terms of two complexes of ideas, one based around the notion of stable electron shells in atoms, and the other around electrostatic forces. These two ideas may be described as explanatory principles (§2.11.2) which acted as the foundation for much of Annie's thinking about chemical bonding. It is possible to interpret much of the development of Annie's understanding about chemical bonding in terms of these two explanatory principles, as over the four interviews the balance of Annie's explanations shifted from being largely based on her stable shells explanatory principle (which was not valid from a curriculum science perspective), to being increasingly constructed in terms of her electrostatic forces explanatory principle (which reflected the principles and explanations of curriculum science).
§7.1: Lability in cognitive structure – evidence from the case study
Annie's progression in understanding chemical bonding may be appreciated by considering some of the conceptual 'tools' (§1.7.2) that she developed or added to her conceptual tool-kit.
Covalent bonding
During the first interview Annie saw covalent bonding as being the type of bonding between two non-metallic atoms ("covalent is a bond that is formed between non-metals", A1.71), whereas she seemed to consider that sharing electrons (or overlap, or combining of atoms) was a more general criterion for the chemical bond (A1.134, 242, 354, 426 and 436).
127 I:… are there any bonds in that diagram do you think?
128 A: Yes.
129 I: How many?
130 A: Four.
131 I: Four bonds, so we've got four bonds there. Erm, are the bonds actually shown?
132 A: Yeah.
133 I: So how are they represented on the diagram?
134 A: By the circles that overlap …
A1
By the second interview her idea of covalent bonding was related to the sharing of electrons between similar atoms (A2.401), although she had little appreciation of the electrostatic nature of the bond. At the start of the fourth interview however Annie was also able to explain how the atomic nuclei attract the bonding electrons due to electrostatic force (A4.10).
1 I: perhaps before we look at any pictures you might just tell me what you think a chemical bond is:
2 A: Chemical bond, erm, it's a link between two atoms, which can be of a various, various different types. But basically links two things together, by either combination, or just by, charge. … just by force they're held together. Actual forces on the atoms. …
10 A: 'Cause the electrons are sort of held in circuits, orbitals, because when they sort of combine together, they're sort of going around freely, so you've got all the forces, sort of just like they're being pulled in by the nucleus. Electrons are being pulled in, so you're, you've got sort of the nucleus pulling in, the electrons from the other, atom. So it helps them stay together.
11 I: What kind of forces are they?
12 A: Electrostatic.
A4
Metallic bonding
During the first interview Annie did not believe metals needed any bonding to hold together, as the atoms involved were of the same element (A1.297).
294 I: Do you think those atoms will hold together?
295 A: Yes.
296I: Why do you think that is?
297 A: Because they're all the same sort.
298 I: Does that make them hold together?
299 A: Yeah.
300I:
Yeah? Do you think there is any kind of bonds between the atoms?
• • • • • • • • • [pause, c.9s]
301 A: No, because they're all the same and they don't need to be bonded.
A1
By the third interview she agreed there was a form of bonding, but as this did not involve atoms combining she seemed to rate this as a lesser form of bonding than covalent (A3.408, c.f. §11.7.2). In the final interview, although Annie still did not consider iron had "actual bonds" she was able to give a description of the "delocalised" electrons which were "like a sea" (A4.82).
81 I: What holds a metal together?
• • • • • • • • • [pause, c.9s]
82 A: Erm, you haven't got like actual bonds in metallic bonding, like you haven't got anything, literally going in or out of a, a, a metal, but you've got delocalised electrons going round, the metallic atoms. In a sort of like a sea. So they're, they're all sort of freely flowing around.
83 I: Why should that hold it together?
84 A: Because, sort of, erm, metals haven't got full, full outer shells, then by electrons moving around, they're, they're getting, er a full outer shell, but then they're sort of losing it, but then like the next one along will be receiving a full outer shell. So, you've also got charges, that are forces from the nucleus pulling, just attracting, atoms from out, or electrons from outside in. Erm,
• • • [pause, c.3s]
A: but mainly due to, like delocalised electrons they can move about, so, then you've got forces keeping, keeping it all together.
A4
Canonical forms
Although in the second interview Annie thought that the canonical forms meant to represent molecular resonance implied discrete molecular structures (A2.299, c.f. §9.4), by the third interview she realised that they were just pictures that were meant to imply delocalisation, and that these chemical structures existed only in the minds of scientists (A3.269)
"Yeah obviously the, • • • [pause, c.3s] sort of all the, all the carbons are going to be sort of have bonding power of four anyway. But sort of where they are actually bonded. It won't affect the structure or the way, in which sort of the, the compound reacts. But it just shows where the bonds could lie, but whether, they don't really exist, it's sort of something that scientist has in their minds to show, to explain something. So sort of three out of the six could be in one position or they could be, in the sort of reverse, although, sort of, I don't know if I should say in nature, they don't actually perform that way."
A3.269
Dative (co-ordinate) bonding
In the first interview Annie demonstrated no concept of dative bonding (A1.626), but in the second interview her ideas of bonding had become sophisticated enough for her to suggest that in some bon
both electrons come from the same atom – even though she used her own nomenclature for this (chlorine bonds, apparently by analogy with hydrogen bonds), rather than the accepted terms dative or coordinate (A2.368),
"[chlorine]'s got a valency [sic] of seven. So, like it would need one electron, so some of the bonds, between like the aluminium and the chlorine, say one out of the four, may, might actually be like a chlorine- chlorine bond, but as the like electrons move round in a circuit anyway you wouldn't be able to trace them. So you can't definitely say."
A2.368
Hydrogen bonding. Annie's comments about a diagram showing a chain of hydrogen fluoride molecules (focal figure 11) illustrate how her ideas on bonding became increasingly sophisticated during her A level course.
focal figure 11
At the time of the first interview she did not recognise the existence of any bonding between the molecules (A1.426).
421 I: Right, okay, do you want to have a look at … picture 11,
A: yeah,
I: figure 11.
A: Oh gosh!
423I: Any ideas about that?
A: {Laughs}
• • • • • • • [pause, c.7s]
A: Not really, but must be a, sort of chain of, hydrogen fluoride molecules.
425 I: Chain of hydrogen fluoride molecules, okay yeah. Erm, is there any kind of bonding going on there?
426 A: There is within the, within the sort of shape of the H-F,
I: uh hm,
A: but when it meets up to like the H-F on the corners of the other shapes, they don't actually bond.
427 I: Okay, how many, how many different H-F molecules can you see there?
428 A: Five.
429 I:Five. And so how many chemical bonds are there in that diagram?
430 A: Five.
A1
In the second interview Annie was able to identify and locate the bond, and comment that it was a lot weaker than a proper bond (A2.268, c.f. §11.7.5),
265 I: Okay, you have a look at number 11? Any idea what type of bonds might be present there?
266 A: Hydrogen bonds.
267 I: Can you tell me where the hydrogen bonds are?
268 A: They're between where the, erm, on the diagram you've got like a, I don't know it's almost like a golf club shape
I: mm
A: where say the, the foot of the club hits the top of the other one, so, if you have H-F, then the, the next one along, the H and the sort of holds them together. Or when you've got the proper bond of H- F, so the actual bond between the, the H and the F of the like neighbouring molecule, is a lot weaker, than the bond, actually in the substance.
269 I: Okay, you say this is a proper bond, this one, between the fluorine and the hydrogen
A: yeah
I: in here. What kind of bond is that?
270A: It's erm, it's a, covalent bond …
A2
It should be pointed out that Annie also wished to locate hydrogen bonds in several inappropriate contexts (A2.6, 59 and 81). This latter tendency seemed to have been overcome by the third interview as by then she was clear that hydrogen bonding could not occur in materials that did not contain hydrogen (A3.82). In addition she was able to explain that this type of bonding was an interaction between a hydrogen atom and a lone-pair of electrons on another atom (A3.84),
"… obviously there's no hydrogen bonding involved [in focal figure 5], 'cause there's not any hydrogen there. Err, probably van der Waals forces, holding. … Sort of van der Waals forces can occur, erm, and, I don't know how to put this. Sort of Van der Waals forces can occur to hold a molecule or atoms together as well as [sic] being sort of involved in bonding, whereas you know, if I was to say like it's hydrogen bonding then that's, involved in just like basically holding molecules near each other like in water the oxygen, lone pairs will attract, to the other hydrogen."
A3.82
Van der Waals' forces
In the first interview Annie seemed to have no concept of van der Waals' forces, and instead invoked alternative, apparently ad hoc, reasons for molecular solids to hold together (A1.730),
727 I: Do you think that … lump of iodine, would it stay together, or do you think it would fall apart?
728 A: Stay together.
729I: Is there something that actually holds it together?
730A: Probably just the forces of pressure and, the, like the charges from each thing would be stable, so…
731 I: What the charges from, what?
732 A: From each molecule.
733 I: So each molecule, is stable,
A: yeah,
and you think that's what is holding it together?
734 A: Yeah.
735 I: Is there any force going from one molecule to another molecule?
736 A: There should be from the forces, the forces from each iodine should have combined to stable-up. But, there's probably other forces, which, erm, hold it together, in a solid or, so it wouldn't, wouldn't break off or anything.
737 I: Right, so there's forces holding the solid together,
A: yeah,
but would they be chemical bonds these forces, or?
738 A: No.
739 I: Not actual chemical, but some other type of force?
740 A: Yeah.
A1
By the later interviews she was clear that van der Waals' forces existed, and that they were weak interactions that were readily disrupted – although she imbued them with an ubiquitous nature and seemed to feel this was a 'catch-all' category that could be applied in a range of inappropriate contexts (A2.2, 93 and 125, A3.82 and 132). So the sodium atom was held together by "van der Waals forces … weak forces, which pull towards the nucleus. Which are readily disrupted" (A2.2). In metallic iron (figure 6) "it's probably van der Waals forces, holding it together" (A2.93), although these forces are not the same as metallic bonding because "you can get van der Waals forces in, covalent things as well" (A2.107). Indeed Annie reported that lithium iodide (figure 8) is "ionically bonded, but the forces holding it together will be van der Waals I suppose" (A2.125).
§7.2: Stability in Annie’s ideas of chemical bonding
The examples given above all show how Annie made significant advances in understanding aspects of chemical bonding during the sixteen months of the case study. However, there were aspects of Annie's thinking where progression appeared to be impeded by the stability of parts of her cognitive structure.
§7.2.1: An example of how G.C.S.E. knowledge can interfere with A level learning: Bond polarity
At G.C.S.E. level students are taught that the chemical elements may be conveniently classed as metals or non-metals (with a few 'semi-metals' or metalloids perhaps mentioned), and this dichotomy amongst elements leads to a dichotomous classification of bonding in compounds – covalent between non-metallic elements, and ionic between a metal and a non-metal (c.f. §11.6).
At A level both dichotomies give way to continua. The elements may be categorised on an electronegativity scale, and bonding may be polar. Essentially covalent compounds may exhibit some degree of ionic behaviour when there is a difference in electronegativity between the elements. Ions may be polarised and 'essentially ionic' substances can show some degree of covalent character.
Annie had clearly learnt that bonding between non-metals is covalent, and between a metal and non-metal is ionic (A1.71, 578, and 744). During her course Annie acquired a concept of bond polarity, which she correctly related to electronegativity (A4.162), and she was also able to discuss the use of the '∂+' and '∂-' symbols to indicate bond polarity, although she was not able to relate this to the notion of partial charges (A3.347). Despite this Annie continued to classify bonding as covalent or ionic, rather than polar ("that'll be, er, ionic … for a start you've got, er metal and a non-metal. And you're going to get complete transfer, of electrons from the lithium to the, iodine atom" A4.284, c.f. §11.6.2). Her G.C.S.E. level knowledge appeared to act as an epistemological learning impediment (§1.5.5) to Annie's progression in thinking about bond polarity.
§7.2.2: An example of how an alternative conception can interfere with orthodox understanding: deviation charges
In the first interview it became clear that Annie's interpretation of the symbols '+' and '-' (which are extensively used in chemistry to show ions) was different to the conventional interpretation. The orthodox meaning is of electrostatic charges, so that any species shown as '+' or '-' is not neutral. Annie however had a totally different interpretation: that the symbols represented deviations from noble gas electronic configurations (A1.262).
Her interpretation led to her not recognising the presence of bonding in a diagram of sodium chloride, as the charge symbols implied to her that the species still had their atomic electronic configurations (A1.238, A3.30).
focal figure 5
One consequence of this was that although Annie interpreted the force between the sodium and chloride species as due to an attraction between opposite charges, for her this meant oppositely signed deviations from noble gas electronic configurations: Na+ being one electron in excess, and Cl- being one electron deficient. As Annie's scheme included 'opposite' charges, and they still attracted, she was presumably still able to make sense of much that she heard and read, despite her alternative interpretation,
256 A: By forces. Any idea what kind of forces would hold it together?
258 A: Probably just the attraction.
259 I: Uh hm.
260 A: 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.
261 I: 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,
I: uh huh,
and chlorine has seven electrons in its outer shell so it's minus an electron so by sort of exchanging,
I: huh hm,
the sodium combining with the chlorine just by force pulls they would hold together.
263I: You say by exchanging, did you say?
264 A: Yeah by, well just the attraction in them.
A1
[focal figure 5] would probably get held together by just forces.
However, this alternative conception did have consequences for Annie's understanding of aspects of her course. One example is that although Annie acquired a reasonably orthodox understanding of the ∂+ and ∂- symbols used to show bond polarity, she did not associate the term 'partial charge' with this symbolism, apparently unable to relate this to electronic configurations (A3.330). Annie was able to 'balance' equations using her deviation charges, but as she was seeking full shells rather than neutrality the results could be quite different to the curriculum science answers: in the case of aluminium sulphate her stoichiometry was (Al3+)4(SO42-)2 rather than (Al3+)2(SO42-)3 (A2.226). Another consequence was that Annie was unsure whether Na+Cl– represented a compound or a mixture of elements, and confused the properties of sodium chloride, with those of its constituent elements (A3.174 and 192).
It is not possible from the case study to suggest the origin of Annie's alternative conception of charge. However it is clear that the alternative deviation interpretation was present in the first interview, whereas there was no evidence of the conventional 'non-neutral' interpretation. By the second interview (after formal teaching of the bonding topic) Annie had acquired the conventional interpretation, but this did not lead to the elimination of the deviation meaning. Indeed her alternative conception appeared to be applied spontaneously, whereas the conventional interpretation was used when questioning was targeted specifically at the electron configuration of ions compared to the atoms. Such cuing appeared to 'switch' Annie into applying her new conventional interpretation, although later she would resort to the alternative meaning.
Annie had presumably made sense of much that had been presented to her at G.C.S.E. and the start of her A level course using a complex of ideas constructed around deviation charges. Revisiting ionic bonding at A level and being taught contradictory ideas must have been confusing, so perhaps it is not surprising that some of Annie's utterances seemed to contain strands of both interpretations . By the fourth interview (after a 'tutorial' intervention) Annie was able to give a good account of ionic bonding in conventional terms, and to apply the conventional application of charge symbols. However, even at this stage there are vestiges of her earlier scheme apparent in the language used, such as referring to a chlorine atom as being "sort of minus an electron" and sodium being a "sort of positively charged, ion because of the, the extra electron" (A4.22 and 26). Like her dichotomy of bond types, Annie's notion of deviation charges delayed her progress, and thus acted as a substantial learning impediment (§1.5.3). (To be precise this would be another example of an epistemological learning impediment as it is "an aspect of cognitive
structure derived from deliberate formal instruction" (§1.5.5). Presumably Annie was never taught the deviation interpretation of charge, but rather – in ignorance of electrostatic ideas – constructed a meaning to interpret her teachers' talk of positive and negative charges. However, a notion about deviations from noble gas electronic structures does not seem a likely intuitive idea, and is not sensibly classed as an ontological learning impediment. As explained in chapter 1, the discrimination between these two categories is intended to inform pedagogic practice, rather than make an absolute distinction. In this particular case it is perhaps most significant that at the time Annie first heard about atoms being ionised she did not have the appropriate prerequisite knowledge about electrical charge: that is she suffered an deficiency learning impediment (§1.5.2), and as in fig 2.3(c) (§2.10.4) thus constructed her meaning in isolation from the curriculum science idea of charge.)
§7.2.3: An example of how the absence of assumed prerequisite knowledge can impede progression
These two examples of stability – unhelpful stability from the point of view of on- going learning – were not the only ones that could have been drawn from the case study.
Another theme that could be explored was her interpretation of diagrams meant to represent electron clouds showing where electron density is most significant in bonds. Annie's understanding – or misunderstanding here – is related to her thinking in other areas. Because Annie did not understand what the diagrams were meant to show they did not help her appreciate polar bonding when it was illustrated through such representations. The reason Annie could maintain an alternative interpretation of the electron clouds as being a type of force–field (e.g. A1.305) was related to her ignorance of basic electrostatic ideas (she did not study A level physics, and had a deficiency learning impediment (§1.5.2) in terms of expected prerequisite knowledge) that had her confuse the effects of charge – distorted electron clouds – with the fields themselves. (This is the same ignorance of fundamental physics that enabled her to believe that neutral atoms would attract if their had opposite deviation charges, whilst remaining skeptical of the attraction between species with orthodox charges.)
§7.3: Two explanatory principles used by Annie to make sense of chemical bond
A consideration of Annie's comments during the research interviews suggests that much of her thinking can be related to two explanatory principles that she applied in responding to questions about chemical bonding. I have labelled these principles the stable shells explanatory principle, and the electrostatic forces explanatory principle.
The stable shells explanatory principle. This principle could be defined as 'chemical bonding is how atoms acquire stable shells'. During the research interviews many utterances were elicited from Annie which may be interpreted as deriving from this principle (see §7.3.1).
In summary form this complex of ideas may be represented:
Some sort of attraction is needed to hold atoms together: this can be of three forms, (a) the formation of chemical bonds by the joining of atoms to form stable shells through sharing of electrons, (b) the combining of atoms due to their [deviation] charges, that is the extra electrons, or the need for additional electrons to form stable shells, or (c) by other forces.
(a) Chemical bonds are also called covalent bonds, and occur between non-metal atoms.
(b) Atoms with matching [deviation] charges, that is metal atoms with non-metal atoms, may combine to form neutral molecules, and this is sometimes called ionic bonding.
(c) The other forces will hold together atoms or molecules that are already stable and have no need to form bonds, or similarly charged metal atoms that are unable to achieve stable shells through combining or joining together. These forces are not proper bonds, and have various names such as metallic bonding, hydrogen bonding, and van der Waals forces.
The electrostatic forces explanatory principle. This principle could be defined as 'chemical bonding is due to the electrostatic force between nuclei and electrons'. This perspective is closely related to the curriculum science model. During the research interviews many utterances were elicited from Annie which may be interpreted as deriving from this principle (see §7.3.2).
In summary form this complex of ideas may be represented:
Opposite charges (positive and negative) attract due to electrostatic force. Similar charges repel.
In an atom the positive charge in the nucleus leads to an electrostatic force which draws the electrons in, and holds the atom together. The strength of the force depends on how close the electrons are to the nucleus. The outer shell electron can not get too close as they are repelled by inner shell electrons.
If atoms collide the charged sub-atomic particles will give rise to forces. Protons in one atom would repel protons in the other, and the electrons would also repel. The protons in one atom would attract the electrons from the other. This may lead to the atoms being held together, with electrons being pulled towards both nuclei, to give a molecule. The force may be greater from one nucleus, and sometimes an electron may be transferred to give ions, which will be held together by electrostatic force in ionic bonding.
In metallic bonding there is a force from nuclei to the sea of electrons; in hydrogen bonding lone pairs of electrons attract hydrogen; van der Waals forces are due to the attraction between opposite charges; in solvation ions are attracted to different parts of the solvent molecule.
These two complexes may be illustrated in terms of the utterances elicited from Annie during the interviews. The notions presented (as the italicised sections §7.3.1 and §7.3.2) are – where not verbatim – in phrases close to Annie's own words. These composites are drawn from all four interviews, so that
(i) the full set of ideas in the complex were not elicited from Annie at any one time;
(ii) during any one interview Annie presented aspects of both complexes.
Citations are provided to the location of the utterances on which these complexes are based, within the interview transcripts. Although there is some evidence of both explanatory principles being used throughout the time she was participating in the research, it is also clear that in the first interview Annie's explanations of bonding were heavily based in the stable shells explanatory principle, with the electrostatic forces explanatory principle only being invoked to prevent atoms falling apart. Over the four interviews this balance shifted so that by the last interview Annie gave for the research the electrostatic forces explanatory principle was much more in evidence, than the stable shells explanatory principle.
§7.3.1: Notions elicited from Annie, related to the stable shells explanatory principle
Some sort of attraction is needed to hold atoms together, either the formation of chemical bonds, or the combining of atoms due to their charges. Bonding involves the joining of atoms which combine to form stable shells. Bonds are represented by circles that overlap (A1.134), and a diagram which does not show overlap does not represent bonding ( A1.426, A1.436, A1.438). For example a diagram such as focal figure 5 which has N a+ and Cl- just in rows, just shows atoms (A1.238) with no bonding (A1.240).
focal figure 5
focal figure 11
Similarly there is no bonding between the molecules in the chain of HF molecules shown in focal figure 11 (A1.446). Chemical bonding involves the sharing of electrons, so the atoms have got two electrons between them, and they have each contributed one to the shell (A1.65), which is called a covalent bond (A1.69). Examples of this include the iodine molecule which holds together because of the sharing of electrons (A2.401); lithium combining with iodine to make a stable outer shell between the two atoms, by sharing electrons (A1.321); the bonds in tetrachloromethane which are covalent as the atoms share electrons to give them all full outer shells (A2.12); hydrogen atoms which combine to form a stable first shell (A1.59); and the bonds in the oxygen molecule which are covalent (A1.226) as each oxygen atom is giving two electrons so they can each form a shell of eight (A1.230).
By comparison, 'just combining' involves a matching-up of deviations from stable shells, that is when something with a positive charge (excess electrons) combines with something with a negative charge (deficient in electrons) to become neutral overall . For example Ca2+ and O2- would just combine because one is lacking two electrons and one has got two, and oxygen, which has a 2- charge, combines with two hydrogens which have a combined charge of 2+ (A3.227). In this context stability relates to electronic structure. For the first shell stable means two electrons, as when two hydrogen atoms are joined because they only have one electron in their first shell , so they combine to form a stable first shell (A1.59). For the second shell stable means eight electrons. Examples of this are oxygen atoms (in oxygen) which give two electrons to 'match-up' so they can form a shell of eight (A1.230), and in sodium when an electron is removed, and the eight electrons in the next quantum shell make the atom more stable than when it had one electron on its own (A4.76). This definition of stability encompasses both elements, so a hydrogen molecule has got two hydrogen atoms, to give an outer shell (A2.135), and compounds, where two or more different elements make up the full stable shell (A2.137).
A full shell implies neutrality, so eight extra electrons would be 8+, which would become nought (A2.236), and would be a neutral charge (A2.234). Similarly if sodium and chlorine were bonded there would be an overall neutral charge, because of donation of electrons neither would then have a plus or minus charge. Other configurations are not stable, so for example hydrogen atoms are unstable because they've only got one electron (A1.77). Not having a stable shell has consequences. It gives rise to charges (which are deviations from a stable shell): hydrogen atoms are minus an electron (A3.150), although two hydrogen atoms would both have a plus charge as they have both got one electron in their outer shell (A4.162). Chlorine is minus an electron, where sodium is over an electron (A1.260), and is positively charged because of the extra electron (A4.26). Iodine has seven electrons in its outer shell so it has a negative charge (A2.109), and carbon atoms have a 4+ charge (A3.273). The species SO42- has two electrons missing off it (A2.170), that is, it is two electrons short (A2.176).
Unstable electronic configurations give rise to forces, so two hydrogen atoms in a molecule would be held together by forces due to their lack of electrons and abundance of them (A1.307), and sodium and chlorine atoms are held together by the attraction from being one electron over, and one short (A1.279). This leads to electron transfer, so sodium loses its extra electron to gain a stable shell (A4.46). Unstable configurations also give rise to 'needs', as in the case of the hydrogen atom, which needs to combine so it can be more stable (A1.77). Species with stable outer shells may be held together, but by other forces, not by chemical bonds. For example, iodine molecules are held together (A1.728) though not by chemical bonds (A1.738) but by other forces (A1.736). Bonding is not needed to to hold the structure of sodium chloride together, just forces (A1.256), and Ca2+ and O2- would not need to form a bond (A1.754), but would just combine (A1.758) to make up full shells (A1.760).
In a metal there are no actual bonds (A4.28), but the structure is held together: the atoms are not really sharing, and are not really combining, but they are held together, so there is something going on (A3.426), and although the atoms are similarly charged they do not repel each other (A3.404). For example there is no bonding in a piece of iron (A2.55), but it is held together by something, probably van der Waals' forces (A2.39). There are forces, as – due to delocalised electrons – an atom is getting a full outer shell , then losing it (A4.84), but this is not as definite as when electrons are completely transferred or shared, so there are not bonds in the sense of covalent or ionic (A4.90).
Electrons are held in place in the atom, which is connected to the set pattern of how many can go in each shell (A1.33).
§7.3.2: Notions elicited from Annie related to the electrostatic forces explanatory principle
Opposite charges attract (A3.42), so where you have got positive and negative they are going to attract (A4.409). This is involved in holding an atom together: electrons do not fall out of atoms, due to the attraction from the protons (A1.43). The attraction is between the electron and the nucleus (A4.201), as protons and electrons have charge (A4.189). The protons in the nucleus have a plus charge (A3.8), and the protons in the nucleus draw the electrons in by electrostatic forces (A3.8). These electrostatic forces come out from the nucleus (A3.6). The larger the atom, the less power the nucleus has on the electrons (A3.8), whereas the closer the electrons are to the nucleus, the more force holds them in (A3.8).
The attraction of opposite charges can also lead to interactions between atoms, for example if two hydrogen atoms collide the proton from one atom could attract the electron from the other atom (A4.205). The attraction can hold atoms to one another, in molecules (A3.146), and remains even when the substance is vaporised(A3.150).This attraction occurs as the nucleus of one atom pulls in the electrons from the other atom (A4.10), with electrostatic force (A4.12), which may lead to a symmetrical arrangement, if the atoms have similar electronegativities (A4.162) or not, in which case ∂ symbols are applied. So in the hydrogen molecule, both nuclei are equally attracted to both electrons (A4.239), so the way the charge has been distributed around the molecule is fairly symmetrical (A4.162), as the distribution of the charge around the molecule has not been polarised (A4.162). H owever, the electrons in an O-H bond would be pulled towards the oxygen more than the hydrogen (A3.306), and so hydrogen would be ∂+ and oxygen would be ∂- (A3.347).
In metallic bonding there are delocalised electrons, like a sea, and there are forces from the nucleus pulling the electrons (A4.84). H ydrogen bonding is involved in holding molecules near each other (A3.84), as lone pairs of electrons attract to hydrogens, so in water oxygen lone pairs attract hydrogen in other molecules (A3.84). Van der Waals' forces occur to hold molecules or atoms together (A3.84), and are due to the attraction of opposite charges (A3.82).
The attraction of opposite charges can lead to the formation of ions, which are atoms that have become charged. '+' represents the electron that has been lost, giving the atom a positive charge, so that K+ is an ion ( A2.143), a potassium atom that has lost an electron ( A2.141), and in N a+ the electron has somehow been removed (A3.64). Ions may be formed by electron transfer, where the nucleus of one atom has the power to draw electrons from another atom in (A4.30), where the force on the electron, is dragging it towards the nucleus (A4.30), as when an electron from sodium is pulled towards the chlorine. The attraction can hold ions to oneanother, which is called ionic bonding (A3.184), as in the chemical bonding between sodium and chlorine (A4.182). This is due to electrostatic force (A4.318), so for example there is an electrostatic force from a potassium ion to a fluoride ion (A4.316), and an electrostatic force from the fluoride ion to the potassium ion (A4.320).
The attraction between charges can lead to solvation. A polar solvent will solvate ions, when the positive ions go to one part of the solvent molecule, and the negative ions to another ( A3.221).
Similar charges repel (A4.517). This has consequences for atomic structure, as it prevents outer shell electrons getting too close to the core (A4.517). This also leads to interactions between atoms, so if two hydrogen atoms collided the two protons, and the two electrons, would repel each other (A4.207).
§7.4: The case in relation to the main themes of the research findings
Annie's case may be used to illustrate major themes which emerged during this research project, and which are illustrated further in chapters 9, 10 and 11.
Annie's progression depended upon the adoption of Coulombic electrostatics as an explanatory principle for bonding
An example of how Annie's developing understanding of chemical bonding depended upon her adoption of conventional electrostatics was her acquisition of the concept of 'hydrogen bond'. Annie progressed from ignorance of this category of bond (§7.3.1), to awareness, to being able to explain it in the case of water as due to an attraction between a lone pair on oxygen, and the hydrogen in another molecule (§7.3.2). This progress may be understood to be related to the adoption of the electrostatic forces explanatory principle as the basis for understanding chemical bonding.
focal figure 5
Another example would be Annie's understanding of the ionic bond. Her initial interpretation of focal figure 5 (a cross section of an NaCl ionic lattice) was that there was no bonding present, and the Na+ and Cl- species were atoms, held together "just by force pulls" (A1.262). Although Annie did not completely adopt an electrostatic model of ionic bonding, by the end of the course she recognised ions, and explained that in the case of potassium and bromine, the potassium would become an ion when it has "got rid of an electron" (A4.405), and there would be "a bromine minus ion from gaining an electron" (A4.407), so "because you've got positive-negative … they're going to attract" (A4.409). In the next chapter it will be shown that colearner Tajinder's developing understanding of chemical bonding also involved his adoption of an electrostatic explanatory principle (§8.4.3).
Annie experienced difficulty in appreciating aspects of the 'orbital' concept used in chemistry
Annie used the terms 'shell' (and 'quantum shell'), 'orbitals' and 'energy levels', but did not seem to clearly discriminate between them (see §9.2.1). Similar problems were experienced by other colearners (§9.2). Annie originally understood the electron density envelopes drawn to represent molecular orbitals as a type of force- field (c.f. §9.2.1).
Annie exhibited beliefs about the interactions of charged particles which are inconsistent with Coulombic electrostatics
Annie's notion of (what I have labelled) deviation charges was an extreme example of alternative notions about electrostatics, as she actually considered electronic configurations themselves to give rise to a force. Annie's misunderstanding of the meaning of the symbols '+' and '-' impeded her progress throughout the course, and at least vestiges of this way of thinking were present in the fourth interview, shortly before her final examination, despite a tutorial intervention after the third interview (§7.2.2).
Annie also spoke of forces in ways that were not in keeping with the curriculum science approach. She demonstrated a belief in a nucleus giving rise to a certain amount of pulling power, rather than construe the force as due to the interaction between charged particles (A4.514, see §10.5.1). Similar findings from other learners will be discussed in chapter 10 (§10.5). So Annie did not understand the reciprocal nature of electrostatic forces, and suggested that when a potassium ion was adjacent to a fluoride ion, the fluoride exerted a larger force on the potassium than vice versa (A4.245, see appendix 31, §A31.4.12). Annie thought a nucleus attracted an electron more than vice versa (and she had a similar perception of the gravitational interaction between the sun and earth). Similar findings from other learners will be discussed in chapter 10 (§10.4.4).
Annie exhibited beliefs constructed from an explanatory principle derived from the octet rule
Annie's stable shells explanatory principle was presented above (§7.3), as was the complex of associated notions elicited from Annie during the research (§7.3.1). As this complex demonstrates, the stable shells explanatory principle was Annie's principal starting point for discussing chemical bonding during the first year of her A level course, and remained a significant basis for her explanations through her second year. A similar explanatory principle was elicited from colearner Tajinder as will be discussed in the next chapter (§8.2.1). Indeed, an explanatory principle of this type seemed to be ubiquitous amongst the colearners in this study (§11.0).
Various aspects of Annie's octet thinking will be shown to be reflected in the data collected from other learners. Particular points to note in this respect are:
- the rationale for bonding: atoms share electrons to obtain full outer shells (§11.1.5);
- anthropomorphic language: atoms give, and share, electrons, and need to combine to become more stable (§11.3);
- a molecular interpretation of the ionic bond: although Annie's notion of deviation charges led to her construing figures representing ionic materials as being pre-bonded, she thought that actual sodium chloride contained molecules (§11.4.3). When thinking about the ionic bond Annie seemed to focus on the act of electron transfer, as if that was the bond (§11.4.2).
- a dichotomous classification of bonding: with models of covalent bonding (sharing electrons by overlapping) and ionic bonding (combining by matching up deviation charges) that do not readily admit of intermediate cases (§11.6).
- bonding is distinguished from other forces: so that where an interaction is not conceptualised in terms of obtaining an stable configuration, it is not classed as a bond (§11.7).
