On this page we will explore ways to support students as they make the transition from pre-16 atomic model to the orbtial model met in post-16 studies. 

Moving from 2D to 3D

How we represent orbitals and their relative energies obviously has implications for our students' understanding. Many drawings are limited to the 2D space. Consequently, it easy to think of the orbitals in a linear fashion, like a ladder.

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What misconceptions might this introduce that need to be addressed?

It would be easy to visualise the orbitals as somehow lying 'above' the nucleus. However, the orbitals are orientated around the nucleus at the centre of the atom. In a sense, the orbitals are all 'on top of each other.' Conceptually, this is difficult to understand, as we tend to think of increasing energy in a 2D fashion.

The animation helpfully displays the orbitals all together within an atom. Don't forget to make use of the ideas on the previous page, Developing Familiarity. 

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Moving from certainty to probability

The pre-16 atomic model suggests that electrons exist in a specified energy level at a certain distance away from the nucleus. The concentric circles drawn to represent the energy levels serves to reinforce this idea. Furthermore, many students make the mistake of thinking that the concentric circles represent the path of electrons around the nucleus, much like a planet orbiting the Sun. An underlying, yet unrecognised belief, is that we can know both the velocity and location of electron at the same time. 

Unfortunately, quantum mechanics puts this idea to bed. Due to the Uncertainty Principle, it is not possible to know both the location and velocity at the same time. The act of measuring one of these variables, makes it impossible to know the value of the other variable.

This means that we move from certainty to probability. An orbital is defined as the region in space in which we have a 90% chance of finding the orbital (this can be worked out mathematically). That means that 90 times out of 100, the electron will be 'discovered' within the region of the orbital. 

Look at the two images below of a 1s orbital. 

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What misconceptions could be introduced by image 1?

1. The orbital has a definite boundary and some students may think that the electron cannot be located outside of the sphere.
2. Some students might think of the electron as moving around the surface of the sphere ie. like a 3D orbit

What are the advantages of using image 2?

1. The image affirms the idea of the orbital as being probability based. Each of the dots represents a possible location of the electron.
2. It can be shown that the probability of finding the electron is not uniform throughout the orbital but there is a greater concentration nearer the nucleus.
3. There is no finite boundary to the orbital reminding students that an orbtial is defined as the region within which there is a 90% chance of locating an electron.

Which images should you use in the classroom?

Use both! For very pragmatic reasons, image 1 is much easier to draw and used frequently in the classroom setting. However, it is important that students understand the concept behind the image and recognise its limitations and how it is derived from image 2. 

Note: the probability density for s-orbtials other than the 1s orbital are more complex and we are entering the realm of undergraduate chemistry. 

The report Students' Understanding of Orbitals: A Survey provides some interesting background reading to some of the misconceptions and difficulties students have with the idea of orbitals.