Atomic models

 Models in physics and beyond are used to describe or explain a challenging concept in a more visual or tangible way (e.g. the brain is like a library, electric current is like water flow). The atomic model is the best guess we currently have that explains our encounters with particles everyday.

 In the IB course there are few surprises; the experiments you are required to conduct usually verify the theory. However, when Rutherford saw the results of the alpha particle scattering experiment, he was very surprised. The results were not in line with the theory and led to a new visualisation of the atom. Science has often developed out of surprise results. 


Key Concepts

Nucleus

Thomson model - plum pudding

The Thomson model of the atom had the following key ideas:

  • A solid lump of homogeneous positive charge (the "pudding")
  • Small negative regions located randomly throughout (the "plums")
  • Neutral overall

This model explained the existence of discrete particles of negative charge (electrons) but was proved wrong by Rutherford.

Rutherford model - nuclear atoms

The Rutherford model of the atom had the following key ideas:

  • A small, densely positive nucleus
  • Surrounded by negative electrons
  • Neutral overall

 

This model was established as a result of the Geiger Marsden experiment:

  1. Alpha particles (charge +2) were fired at a thin gold foil. A detecting film was situated all around the experiment chamber.
  2. Almost all of the alpha particles passed through (all should have passed through according to the Thomson model).
  3. Some of the alpha particles were deflected by a small angle.
  4. A very small, but statistically signficant, number of alpha particles (~1 in 1000) were deflected by angles greater than 90 degrees and bounced back. This was evidence for a small positive nucleus, with most of the atom as empty space, as the positive charge of the atom was now known to exist in discrete locations.

http://www.thinkib.net/files/physics/activities/Rutherford.gif

Essentials

Electrons and quantum

Quantum physics

So far we have considered light to behave as a wave. However, light also behaves as discrete quantised particles (photons) in certain experiments. Energy and wavelengths of light are linked mathematically so that specific wavelengths correspond to specific energies:

\(E=hf\) and \(E={hc\over \lambda}\)

  • E = energy of a photon (J)
  • h = Planck's constant (\(6.63 \times 10^{-34}\) m2 kg s-1)
  • f = frequency of light (Hz)
  • c = speed of light (\(3\times 10^8\)ms-1)
  • λ = wavelength of light (m)

Atomic spectra

When a low pressure gas is excited, the light emitted is made of discrete wavelengths (responsible for distinct colours of visible light). These wavelengths can be split by a prism to give a line spectrum known as an emission spectrum.

When white light is absorbed by a gas, the transmitted light has an associated absorption spectrum, opposite to that shown here. To excite an electron from a low level to a high one requires a certain amount of energy. This will only occur if the light shining contains exactly the right amount of wavelength to correspond to this energy.

The line spectra of low pressure gases can be explained if the electrons can only have certain discrete energy levels. These can be represented on an energy level diagram. This is the energy level diagram for hydrogen.

Note that, since the energies involved are very small, the units of energy in this diagram are electron volts (eV). This is the amount of energy gained by an electron accelerated across a potential difference of 1 V

Ionisation is the complete removal of an electron from the atom. This is the same as moving to an orbit where the energy is 0 eV. Remember that electron energy is negative (due to the attraction that exists between the electrons and the positive nucleus) so work must be done to remove electrons. It's like they are in a hole.

Bohr model

We have a lot of evidence that electrons exist in discrete energy levels.

In 1913, Niels Bohr predicted the energy levels of the hydrogen atom by assuming that the angular momentum of orbiting electrons =\(nh\over{2\pi}\)  where n is a whole number (quantum number).

Test Yourself

Use flashcards to practise your recall.


Just for Fun

Check out this πg physics summary.