A model approach to teaching atoms

Describing the structure of atoms relies on models. The models used at 11–14, 14–16 and post-16 become increasingly more complex. Understanding the post-16 models is useful when teaching 14–16 students, to minimise student misconceptions.

At 14–16, we model the atom as a tiny, massive, positively-charged nucleus, surrounded by shells of electrons. This model primarily stems from the work of Hans Geiger, Ernest Marsden, Ernest Rutherford and Niels Bohr. Rutherford inferred the existence of atomic nuclei from the gold foil experiments carried out by Geiger and Marsden. Bohr, in his interpretation of atomic spectra, proposed that electrons exist in ‘orbits’ (now shells) with fixed energy levels, and that electrons can move between the orbits when energy is absorbed or emitted.

Potential issues with the model we use come from students wondering ‘what are the shells made from?’, and the inability of the model to account for the three-dimensional shape of simple molecules.

What you need to know

Why do we stop at calcium when describing electronic structures at 14–16? This is a direct consequence of the simplified model we use. We often say there are a maximum of two electrons in the first shell, and eight in the subsequent shells. At calcium (20 electrons, so 2,8,8,2) we hit the limit of this model, as the electronic structure of scandium (21 electrons) is 2,8,9,2, rather than the expected 2,8,8,3 from the model we use.

The reason for this apparent anomaly becomes clear when we look at the more sophisticated model used for post-16 chemistry. While we usually consider electrons to be particles, they are more accurately described as quantum particles. Electrons have both particle-like and wave-like properties. Using the rules of quantum mechanics, we can derive four numbers which identify the position and energy of electrons in atoms.

These quantum numbers help to describe the electron shells, which have a complex substructure, being made of subshells, which themselves are made of orbitals. Within each orbital, two electrons can exist.

Each shell has one more subshell than the previous shell, and each subshell has two more orbitals then the previous subshell. For example, the first shell is formed of one s-type subshell, made up of one orbital that we call 1s. Two electrons with opposite spin can exist in this orbital. This accounts for the electronic structure of hydrogen (1 or 1s¹) and helium (2 or 1s²).

The second shell is formed of an s-type subshell (one orbital, hence two electrons, called 2s) and a p-type subshell (three orbitals, hence six electrons, called 2p). This accounts for the electronic structures of beryllium (2,1 or 1s²2s¹) through to neon (2,8 or 1s²2s²2p⁶).

Clearing up misconceptions

Students at 14–16 will sometimes ponder the structure of the periodic table, and wonder at the apparent arbitrariness of the layout. Why are there only two elements in the first row, then eight in the second and third? Why are these rows divided into two and six like that?

Knowing something about subshells and orbitals starts to draw back the conceptual curtain. If shell one is made of only one orbital, it makes sense for there to be only two elements. If shell two is made of four orbitals, it makes sense for there to be eight elements. And so on. You can even refer to the blocks by their formal names, from left to right as s, d and p (with f usually underneath for convenience).

As always with models, they are a description of reality at a level useful for a purpose. For the majority of students, understanding electron structures up to calcium is sufficient, so there is no direct need to introduce a more complex model. However, there are some concepts at 14–16 that can’t be explained by the model we use, and so it can be useful to allude to the more sophisticated model.