Welcome to practical physicsPracticle physics - practical activities designed for use in the classroom with 11 to 19 year olds
 

Developing a model of the atom: the nuclear atom

Students will have seen signs of electrons and positive ions (perhaps carrying a current through ionized neon or helium). On this evidence, Thomson proposed his ‘plum pudding’ model, in which the negative electrons sit in the positive nucleus like currents in a bun. 
 
However, students may well have been exposed to images, stories, films and textbooks which incorporate a version of a planetary model. They may find it strange that you even mention the plum pudding model. It is not difficult to persuade them that this model has been superseded. What is more challenging is to take them beyond the simple pictures that they may have seen and to give them some idea of the size of atoms and nuclei – i.e. that the atom is not just hollow but extremely hollow (too hollow to represent in simple images). 
 
The Rutherford-Bohr model 
Rutherford proposed the idea of a central nucleus in an atom in around 1908. The nucleus contains the atom’s positive charge whilst the electrons are outside. From back scattering experiments, he showed that the radius of the nucleus was 100,000 times smaller than the radius of an atom. This is equivalent to the head of a pin (the nucleus) in the middle of a large stadium (the atom). Consequently, the relative sizes of the atom and its nucleus cannot be shown in simple diagrams. 
 
Given that the electrons were easily removed, Rutherford assumed that they were on the edge of the atom. There emerged a picture of orbiting electrons that mirrored the planets orbiting the Sun.
 
Problems with the planetary model 
However, whilst the existence of the hollow atom is well accepted, there have always been serious objections to a classical planetary model. The main objection is that the orbiting electrons are moving charges and should radiate electromagnetic waves, losing energy. This loss of energy would cause them to spiral into the nucleus. In other words, there was no way of explaining why an atom with orbiting electrons is stable. 
 
The Bohr model 
This issue of instability was addressed by Niels Bohr in 1913. He combined Rutherford’s model with the quantum ideas put forward by Max Planck at the turn of the century. Bohr no longer referred to orbits but only to ‘stationary states’. Atoms could exist in a stationary state and be stable. Spectral lines were a result of transitions between these stationary states. Bohr’s model was the first step towards an atom that is described by quantum mechanics. It no longer followed classical laws – particularly those of electrodynamics. The reason for this was straightforward: classical electrodynamics could not explain how the electron could be bound to the nucleus andbe stable. 
 
In the late 1920s, the quantum mechanics of Schrödinger and Heisenberg offered two further developments to Bohr’s model (see The atom - a quantum mechanical model). These models become ever more difficult to represent in simple images. Perhaps because of this, the simple planetary picture has endured as one of the icons of twentieth century atomic physics, even though it was superseded within a few years of being proposed.