Use post-16 atomic models to inform understanding of 14-16 | CPD

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The description of the structure of atoms is based on models. The models used at 11-14, 14-16 and after 16 are becoming increasingly complex. Understanding the post-16 models is helpful when teaching 14-16 students to minimize student misunderstandings.

In 14-16 we model the atom as a tiny, massive, positively charged nucleus surrounded by shells of electrons. This model derives primarily 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 of Geiger and Marsden. Bohr proposed in his interpretation of atomic spectra that electrons exist in “orbits” (now shells) with fixed energy levels, and that electrons can move between orbits as energy is absorbed or emitted.

Potential problems with the model we use come from students wondering, “What are the shells made of?” and the model’s inability to account for the three-dimensional shape of simple molecules.

What the model lacks

Ideas that cannot be explained with the 14–16 model of electronic structure include:

  • Binding in H2SO4
  • why transition metals are strongly colored
  • why zinc is actually not a transition element
  • why oxygen is paramagnetic
  • why he2+ can exist, but He2 tip.

what you need to know

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

The reason for this apparent anomaly becomes clear when we look at the more sophisticated model used for post-16 chemistry. While we usually think of electrons as 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 that identify the position and energy of electrons in atoms.

These quantum numbers help describe that the electron shells, which have a complex substructure, are made up of subshells, which in turn are made up of orbitals. Two electrons can exist in each orbital.

Each shell has one more subshell than the previous shell, and each subshell has two more orbitals than the previous subshell. For example, the first shell consists of a type S subshell made up of an orbital we call 1s. Two electrons with opposite spin can exist in this orbital. This explains the electronic structure of hydrogen (1 or 1s1) and helium (2 or 1s2).

The second shell consists of an s-type (one orbital, hence two electrons, called 2s) subshell and a p-type (three orbitals, therefore six electrons, called 2p) subshell. This explains the electronic structures of beryllium (2,1 or 1s22s1) to neon (2.8 or 1s22s22p6).

The third shell consists of one 3s orbital, three 3p orbitals, and five 3d orbitals. There are complex rules that explain the order of filling the orbitals. Since the electrons fill the 3s and 3p orbitals, we declare sodium to argon. However, in potassium and calcium atoms, the 3d orbitals have higher energy than the 4d orbitals, so the latter orbitals fill with electrons first. In scandium, the relative energies of the 3d and 4s orbitals are reversed, and one of its 21 electrons goes into a 3d orbital. The particular shapes of the orbitals shown in the figure above directly affect the shapes and properties of molecules. For example, s orbitals are generally spherical and p orbitals look a bit like dumbbells. How these orbitals overlap when electrons are shared in covalent bonds determines the strength and type of covalent bonds formed.

I teach my 14-16 year old students and will not go into the full details of subshells and orbitals (unless they are particularly interested). However, I will specifically address the limitations of the model we use and why we make seemingly arbitrary decisions about which electronic structures of the elements interest us. Keith Tabers An analogy for the atom probe is a useful resource if you want to continue these discussions.

The third shell consists of one 3s orbital, three 3p orbitals, and five 3d orbitals. There are complex rules that explain the order of filling the orbitals. Since the electrons fill the 3s and 3p orbitals, we declare sodium to argon. However, in K and Ca atoms, the 3d orbitals are higher in energy than the 4s, so the latter orbitals fill with electrons first. In scandium, the relative energies of the 3d and 4s orbitals are reversed, and one of its 21 electrons goes into a 3d orbital. The special shapes of the orbitals, shown in the figure below, directly affect the shapes and properties of molecules. For example, s orbitals are generally spherical and p orbitals look a bit like dumbbells. How these orbitals overlap when electrons are shared in covalent bonds determines the strength and type of covalent bonds formed.

I teach my 14-16 year old students and will not go into the full details of subshells and orbitals (unless they are particularly interested). However, I will specifically address the limitations of the model we use and why we make seemingly arbitrary decisions about which electronic structures of the elements interest us. Keith Tabers An analogy for the atom (rsc.li/3PIL6fk) probe is a useful resource if you want to continue these discussions.

clear up misunderstandings

Students aged 14 to 16 sometimes ponder the structure of the periodic table and marvel at the apparent randomness of the layout. Why are there only two elements in the first row, and eight in the second and third? Why are these rows so divided into two and six?

Knowing about subshells and orbitals begins to pull back the conceptual curtain. If shell one consists of only one orbital, it makes sense that there are only two elements. If shell two consists of four orbitals, it makes sense that there are eight elements. Etc. You can even refer to the blocks by their formal names, from left to right as s, d, and p (usually with f below for convenience).

As always with models, they are a description of reality at a level that is useful for a specific purpose. For the majority of students, understanding the electronic structures down 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 cannot be explained by the model we are using, and so it may be useful to point out the more sophisticated model.

Suggestions for teaching atomic models at 14-16

  • Tell your students openly that 2,8,8,2 is a model, a rule of thumb – useful for explaining some ideas but showing only some of the superficial details of what is really happening.
  • Ask your students how they visualize atoms. Do you see them as physical shells in which the electrons sit?
  • Try to be careful and consistent with your language. Avoid jumping back and forth between “energy level” and “electron shell” when discussing these ideas.
  • Use analogies when helping students formulate their concepts. For example, electrons fill shells like hotel guests do rooms in a high-rise hotel. Higher energy shells (higher floors) are empty (available) to allow electrons (guests) to move in when needed.
  • Encourage interested students to continue reading – refer them to general post-16 textbooks (e.g. Hill and Holman’s chemistry in context) or websites such as chemguide.co.uk.
  • Make sure your students can write concise and coherent answers to exam-like questions such as “Describe and explain the electronic structure of a calcium atom.”

Suggestions for teaching atomic models at 14-16

  • Tell your students openly that 2,8,8,2 is a model, a rule of thumb – useful for explaining some ideas but showing only some of the superficial details of what is really happening.
  • Ask your students how they visualize atoms. Do you see them as physical shells in which the electrons sit?
  • Try to be careful and consistent with your language. Avoid jumping back and forth between “energy level” and “electron shell” when discussing these ideas.
  • Use analogies when helping students formulate their concepts. For example, electrons fill shells like hotel guests do rooms in a high-rise hotel. Higher energy shells (higher floors) are empty (available) to allow electrons (guests) to move in when needed.
  • Encourage interested students to continue reading – refer them to general post-16 textbooks (e.g. Hill and Holman’s chemistry in context) or websites such as chemguide.co.uk.
  • Make sure your students can write concise and coherent answers to exam-like questions such as “Describe and explain the electronic structure of a calcium atom.”

materials for your lessons

materials for your lessons

  • Encourage students to explore the key scientists involved in developing our understanding of atomic structure. Start with a biography by Niels Bohr: bit.ly/3w4mYMj
  • Use this classroom activity to quiz your students about the atomic structure of atoms: rsc.li/3zGCpwx
  • Break student misconceptions about chemical stability with these ready-to-use teaching materials: rsc.li/3zBKNNK

David Paterson teaches chemistry at Aldenham School and is a chemistry consultant at CLEAPSS

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