AtomsIntroduction
It is Diwali, and the mela is a blaze of light: red neon signs over the stalls, an orange sodium lamp on the street, and a tawa glowing red-hot at the jalebi stall. Riya has brought a small prism from the school lab. Through it, the tawa's glow spreads into a smooth rainbow. But the street lamp does something strange.
By the end of the nineteenth century there was plenty of evidence that matter is made of atoms. Hot solids and liquids, where the atoms are packed close together and constantly jostle their neighbours, give out light of every wavelength: a continuous spectrum. A thin gas made to glow, in a flame or in a discharge tube like a neon sign, behaves very differently. Split the light of each source and look.
In a rarefied gas the atoms are far apart, so the light comes from the atoms one at a time, not from their collisions with neighbours. Each element gives its own fixed set of bright lines, its line spectrum. Hydrogen always gives the same four visible lines in the same places, sodium always its close yellow pair. The pattern is so dependable that it identifies the gas, like a fingerprint, and it strongly suggests that the light is tied to what goes on inside each atom.
In 1885 Johann Jakob Balmer, a Swiss school teacher, found a simple formula that gave the wavelengths of the hydrogen lines. Nobody could say why it worked. Hydrogen is the simplest atom of all, one electron and one proton, so it became the test case: any model of the atom had to explain its spectrum. This chapter follows the models that tried.
Watermelon or Solar System?
In 1897 J. J. Thomson's experiments on electric discharge through gases showed that atoms of every element contain the same tiny negative particles, the electrons. An atom as a whole is neutral, so it must also hold an equal amount of positive charge. But how are the two arranged?
In Thomson's model of 1898, the positive charge is spread evenly through the whole atom and the electrons sit in it like seeds in a tarbooz (watermelon). It was nicknamed the plum pudding model. Rutherford's model of 1911 puts all the positive charge and nearly all the mass in a tiny nucleus at the centre, with the electrons going round it like planets round the Sun. Nothing in the picture is to scale: the real nucleus would be a speck too small to draw.
Rutherford's nuclear model, built on the α-particle experiments of Section 12.2, is close to how we picture the atom today. But it left one big question open. Which of these could it NOT explain?
A planet-like electron could have any orbit and any energy, so there is no reason for a single electron and a single proton to give only a handful of sharp wavelengths. Worse, as Section 12.4 shows, classical physics says such an atom cannot even be stable. Fixing both problems needed Niels Bohr's quantum ideas.