Topic G: Atomic World (Elective)

The Atomic World elective explores the experiments and ideas that revealed the structure of the atom and the strange dual nature of light and matter. You will follow the path from the nuclear atom to the quantum picture, see how the photoelectric effect forces us to treat light as particles, and learn how we probe matter down to the nanometre scale.

The Rutherford nuclear atom

In the alpha-particle scattering experiment, a thin gold foil was bombarded with alpha particles. Most passed straight through, showing the atom is mostly empty space, but a tiny fraction were deflected through large angles and a very few bounced straight back. Rutherford concluded that the positive charge and almost all the mass of an atom are concentrated in a very small central nucleus, with electrons orbiting at relatively large distances. This nuclear model replaced the earlier plum-pudding model. However, classical physics predicted that orbiting electrons should continuously radiate energy and spiral into the nucleus, so the Rutherford atom could not be stable, pointing to the need for a new quantum description.

The photoelectric effect

When light of high enough frequency shines on a metal surface, electrons are ejected; this is the photoelectric effect. Experiments show that emission depends on the frequency of the light, not its brightness: below a certain threshold frequency no electrons are emitted however intense the light, while above it electrons come out immediately. This cannot be explained by treating light as a continuous wave. Einstein proposed that light comes in packets called photons, each carrying energy E = h f, where h is Planck's constant and f is the frequency. A photon gives all its energy to one electron; if this exceeds the work function (the energy binding the electron to the metal), the electron escapes with the surplus as kinetic energy.

The Bohr model and energy levels

Bohr modified the nuclear atom by proposing that electrons can occupy only certain allowed orbits, each with a fixed, quantized energy, and that an electron in such an orbit does not radiate energy. An electron can jump between levels only by absorbing or emitting a photon whose energy exactly equals the difference between the two levels. This explained the stability of the atom and, crucially, the line spectra of hydrogen. The model introduces the idea that energy on the atomic scale is quantized rather than continuous, a radical departure from classical physics, although the simple Bohr picture works well only for hydrogen and other one-electron systems.

Atomic spectra

Because the energy levels in an atom are quantized, the photons an atom can emit or absorb have only certain energies, and therefore only certain frequencies and wavelengths. When excited atoms drop to lower levels they emit light at these specific wavelengths, producing a line emission spectrum of bright coloured lines on a dark background. When white light passes through a cool gas, the same wavelengths are absorbed, giving a line absorption spectrum of dark lines. Each element has a unique pattern of lines, a fingerprint that allows elements to be identified, in flame tests, in stars and in chemical analysis. The convergence of the lines at high frequency marks the ionisation energy of the atom.

Wave-particle duality

The photoelectric effect shows that light, normally treated as a wave, can behave as a stream of particles, while interference and diffraction show it behaving as a wave. This dual behaviour is called wave-particle duality. De Broglie extended the idea to matter, proposing that particles such as electrons also have an associated wavelength given by wavelength = h over p, where p is the momentum. This was confirmed when beams of electrons were diffracted by crystals, producing patterns just like those of X-rays. The wavelength of fast electrons is very small, which is why electron microscopes can resolve far finer detail than light microscopes. Duality shows that the classical separation of waves and particles breaks down at the quantum scale.

Probing matter

To investigate structures smaller than the wavelength of visible light, scientists use probes with much shorter wavelengths. Electron microscopes accelerate electrons to high speeds so their de Broglie wavelength is tiny, allowing them to resolve viruses, molecules and fine surface detail far beyond the reach of optical microscopes. X-ray diffraction uses the regular spacing of atoms in a crystal as a diffraction grating; the pattern of scattered X-rays reveals the arrangement of the atoms and was used to determine the structure of DNA. Scanning probe microscopes can even image and manipulate individual atoms. In every case the rule is that to see fine detail you need a probe whose wavelength is comparable to or smaller than the structure.

Nanoscience and nanotechnology

Nanoscience studies matter on the scale of nanometres, where one nanometre is one billionth of a metre, roughly the size of a few atoms. At this scale materials can show properties quite different from those of the bulk material, because a much larger fraction of their atoms lie on the surface and quantum effects become important. Examples include carbon nanotubes, which are extraordinarily strong and conduct electricity, and nanoparticles used in catalysts, sunscreens and medicine. Nanotechnology aims to design and build useful devices atom by atom, with applications in electronics, materials and targeted drug delivery, while raising new questions about health and environmental safety that are still being studied.