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New Zealand Science Teacher

Science Curriculum/Scientific Literacy

The nature of chemical bonding

Sheila Woodgate at the University of Auckland, through an analysis of the responses of 400 Year 12 and Year 13 students to a BestChoice module about ionic, covalent and metallic bonding, has exposed student misconceptions that could inform the teaching of bonding.

The module

The first part of the module focuses on the electrostatic nature of bonding and the nature of the positive-negative attractions for the three bond types. An introductory information page points out that bond formation is exothermic, and as a consequence, the energy of a system in which atoms are joined by chemical bonds is lower than the energy of the corresponding non-bonded atoms. It also describes that bonding models attribute the lowering of energy for all types of bonds to a larger number of attractive electrostatic (positive-negative) interactions in the bonded system, and the nature of the positive and negative particles for each bonding type.

Despite this information being available, student responses indicate that that they believe the term ‘electrostatic’ applies only to ionic bonding. When asked “Which type of bonding involves attractions between positive and negative particles?” 20% of students chose “covalent, ionic and metallic”. Most students chose “only ionic”. 80% of students could identify the type of positive and negative particles involved in ionic bonding, but only 50% could identify the positive and negative particles attracted to one another in metallic or in covalent bonding.

Issues raised

For covalent and metallic bonding, is it enough to tell students that electrons are shared?

In a description of covalent and metallic bonding, it is important also to mention the positive particles that are sharing the electrons because it is the sharing of electrons by positive particles that lowers the energy of the bonded system. In a covalently bonded system the bonding electrons are between two nuclei and are attracted to (shared by) both of them. In a metal lattice, each metal atom can be thought of as bonding electrons and a cation (Na+ + e-). In the accepted model of metallic bonding the bonding electrons are attracted to (shared by) the cations in the lattice. This may be described as cations in a sea of electrons. It is not correct to describe metallic bonding as electrons shared between atoms. Students could conclude that the shared electrons are additional to those present in atoms. Furthermore, why should there be a strong attractive force between electrons and uncharged atoms?

Do we focus too much on ion formation and not enough on the attractive forces in an ionic compound?

It is common and misleading to associate ionic bonding with ‘transfer of electrons’. Transfer of electrons gives rise to the particles that are attracted to one another. However, it is the bringing of the ions together in the lattice (not the ion formation) that is responsible for release of energy in formation of an ionic compound. Electron transfer is at best slightly exothermic and may be endothermic, even if the ions formed have filled shells. Do not take for granted that students associate the bringing together of particles of opposite charge with release of energy. 485 Year 13 students were asked in BestChoice to identify the lowest energy system from four choices with positive and negative spheres at various degrees of separation, 67% choose the pair with the largest degree of separation!

Is indiscriminate use of the word ‘bond’ potentially misleading?

The word “bond” implies two atoms as in a covalent bond. Ions in ionic solids and metal atoms in metals are bonded to all atoms (ions) that surround them. This is better described as metallic bonding or ionic bonding.

Is our symbolic representation of metals misleading students?

Symbolic representations of elements and compounds are introduced early in studies of chemistry. Teaching of these should be associated with discussion of the form taken by elements to emphasize, for example, that while Ne(g) represents neon atoms, Na(s) does NOT represent a single atom but a collection of atoms in a lattice. BestChoice data also show how students can misinterpret symbols. A question in the bonding module asks them to identify the form (atom/lattice/molecule) of various substances (including Na solid, NaCl solid). 92% of 395 answered that NaCl exists as a lattice, but only 41% recognised that Na exists as a lattice. 33% thought that “Na solid” was an atom, and as a consequence was not involved in bonds. Describing states of aggregation in the elemental form using the terms atom, molecule (a small collection of atoms), and lattice (a large collection of atoms in an ordered arrangement) is as important as distinguishing between an element and compound.

Are the types of bonding really so different?

The bonding types can be placed at the corners of a triangle, a representation that highlights both connections and transitions between them (http://tinyurl.com/7repxmr). One way of further developing understanding of bonding is to point out connections between bonding models. Lessons that describe the transition from covalent to ionic bonding as sharing of electrons becomes unequal could also include comparisons between metallic bonding and covalent or ionic bonding. Both metallic and ionic bonding involve attractive forces between an atom (ion) and all surrounding atoms (ions) in a lattice. Metallic and covalent bonding have in common, the sharing of electrons by positive particles, but differ in the extent of sharing. In metallic bonding the bonding electrons are shared by all metal cations in the lattice, and in covalent bonding electrons are shared by two atoms.

Furthermore, the bonds between atoms in molecules and the attractive forces between molecules are all electrostatic in nature. It makes no sense to ignore or gloss over the electrostatic nature of a covalent bond and to put such an emphasis on the electrostatic nature of weaker attractions between molecules due to permanent and temporary dipoles.

For further information contact: Suzanne.Boniface@vuw.ac.nz

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