ANTECEDENTES DEL ESTUDIO

Intermolecular bonds do not normally feature in pre-16 chemistry courses in the UK. Ideas about hydrogen bonding, other types of dipole-dipole bonds including those frequently termed “van der Waals’ forces” are taught in post-16 courses. The topic has received relatively little attention from chemical education researchers.

10.3.1 Hydrogen bonds

Hydrogen bonds arise when hydrogen is bonded to the highly electronegative elements fluorine, oxygen and nitrogen. For example, in hydrogen fluoride, the electrons in the covalent bond between hydrogen and fluorine are distributed towards the electronegative element, distorting the electron cloud and creating permanent positive and negative charges

Beyond appearances: students’ misconceptions about basic chemical ideas 2 edition 2004

on the molecule, referred to as a “dipole”. The hydrogen nucleus contributes the positive charge and the distorted electron cloud around the fluorine atom takes a negative charge. The positive charge from one molecule may align with the negative charge on another, resulting in a specific type of electrostatic attraction called a “hydrogen bond”.

Progression in the development of basic ideas

Barker (1995) and Taber (1993a) have explored students’ thinking about hydrogen bonds. In Barker’s survey, 250 students beginning post-16 chemistry study were asked to identify the bonds between water molecules and to explain what distinguished these from covalent bonds. At the start, about 18% identified these as hydrogen bonds, increasing to about 69% fifteen months later. About 20% began by suggesting the bonds were “liquid” bonds or “ weak” bonds between molecules, possibly because a lack of formal teaching led to guessing from the diagrams provided. About 8% at the first stage described hydrogen bonds as “an attraction force, not a bond”. Fifteen months later, few students gave the “liquid/weak” bond response, but 24% gave the “attraction” description. This suggests that students learn to distinguish between intermolecular bonds and other types of bond, and ascribe these different properties. This is neither chemically accurate or necessary.

Taber’s work with Annie (1993a) gives a more specific view of progression in understanding of hydrogen bonds. Annie was presented with a diagram representing a chain of hydrogen fluoride molecules. The molecules were shown with the appropriate distorted electron cloud, and were drawn touching one another. Annie did not think any bonding was present between the molecules. Taber suggests this may have been because the shapes did not overlap one another. In her second, post-teaching interview, Annie could describe the difference between the O-H bond within a water molecule and the bond between two water molecules:-

"You've got the two hydrogens added to an oxygen. And then the hydrogen brings a small bonding between like another oxygen, to hold the structure together but it's not like, it is a bond, but it's not as strong, as like, the ionic bond would be" (p 42).

In her third interview, Annie talked about hydrogen bonds involving lone pairs of electrons and demonstrated much clearer understanding of the intermolecular role of hydrogen bonding.

10.3.2 Other intermolecular bonds

Beyond appearances: students’ misconceptions about basic chemical ideas 2 edition 2004

Temporary positive charges bond with temporary negative charges. This type of interaction can be called a “van der Waals’ force”. Each electrostatic attraction is small in energy terms, but when thousands or millions are being made and broken their effect on the structure and function of a substance is significant.

Barker explored students’ thinking about intermolecular bonds other than hydrogen bonds by asking students to explain why the vapour at 1000 oC above a mixture of titanium(IV) and magnesium chlorides comprised titanium(IV) chloride only, given that titanium(IV) chloride is “covalent” and magnesium chloride “ionic” in nature. At the start, only 1% of respondents suggested that intermolecular bonds between titanium(IV) chloride molecules would break, a figure which increased to 16% fifteen months later. Initially, students starting post-16 chemistry study divided into four groups. Those who thought that covalent substances have lower boiling points, so more heat was needed to vapourise the magnesium chloride numbered 22%. About 13% thought that ionic bonds can’t be broken by heating. Almost one-quarter (24%) suggested that covalent bonds are weaker than ionic bonds so break. About one-third (33%) gave no response or an uncodeable response. By the end of the study these responses were still prevalent; the figures giving these answers were 14%, 15% and 31%, with 11% giving an uncodeable or no response. These data point to the widespread use of qualitative and vague ideas focusing on the behaviour of substances, despite the fact that the course followed by these students presented all intermolecular bonds in a chemically correct, context-led way.

At her first interview, Annie (Taber, 1993a) was asked about the structure of iodine. She explained that iodine molecules were held together by "forces of pressure", not chemical bonds. After teaching, she was aware of the existence of van der Waals' forces, and correctly placed these between iodine molecules, but thought that they would also occur in compounds like sodium chloride, as though she was applying them to any structure which she could not otherwise explain. Annie knew at this second stage that van der Waals' forces would be affected by heat, but could not explain this in an accepted way. In her final interview, Annie retained the idea that van der Waals' forces existed in sodium chloride, and realised that these bonds would break before covalent bonds when a substance was heated. Annie’s views support those reported in the large scale study.

Associated difficulties

In learning about intermolecular bonds some students develop misconceptions. One common error touched on by Annie and reported more formally by Peterson and Treagust (1987) is misunderstanding of the different locations of inter- and intramolecular bonds. About 23% of students thought that intermolecular bonds were within a covalent molecule. In his later study, Peterson (1993) found that 36% of first year university chemists thought that silicon carbide had a high melting point because of "strong intermolecular forces".

Beyond appearances: students’ misconceptions about basic chemical ideas 2 edition 2004

Students also misunderstand the relative strengths of inter- and intramolecular bonds. Peterson and Treagust report that one-third of their sample of Australian sixth formers thought that "strong intermolecular forces exist in a continuous covalent network" (p 460).

10.4 Summary of key difficulties