Atoms are so small that millions of them lined up side by side would just span the dot at the end of this sentence. No one could see one for most of history, so how do we know what they look like inside? Scientists built models: simplified pictures or descriptions that explain what we observe. Each time a new experiment produced a result the old model could not explain, the model was changed. The story of the atom is one of the best examples of how science works: ideas are tested, and evidence wins.
The idea of atoms is old. Around 400 BCE the Greek thinker Democritus argued that if you kept cutting matter in half, you would eventually reach a piece that could not be cut. He called it atomos, meaning uncuttable. But he had no experiments to test the idea, and it was mostly set aside for about 2000 years.
| Scientist and year | Model | Key evidence or idea |
|---|---|---|
| John Dalton, 1803 | Solid sphere (billiard ball): atoms are tiny, solid and indivisible | Elements always combine in fixed mass ratios to form compounds |
| J. J. Thomson, 1897 | Plum pudding: electrons scattered through a ball of positive charge | Cathode rays are streams of tiny negative particles, the electrons |
| Ernest Rutherford, 1911 | Nuclear model: a tiny, dense, positive nucleus with electrons around it; mostly empty space | Gold foil experiment by Geiger and Marsden (1909) |
| Niels Bohr, 1913 | Electrons travel in fixed orbits called energy levels or shells | Each element gives off only certain colours of light (a line spectrum) |
| James Chadwick, 1932 | The nucleus contains neutral neutrons as well as protons | A neutral radiation that knocked protons out of wax |
Dalton's model (1803). Chemists had discovered that elements always combine in the same mass ratios: water always contains 8 g of oxygen for every 1 g of hydrogen. The English chemist John Dalton explained this with an atomic theory, first presented in 1803 and published in full in 1808: all matter is made of tiny, indivisible atoms; atoms of one element are identical; atoms of different elements differ in mass; and compounds form when atoms join in simple whole-number ratios. His picture of an atom was a solid ball, like a tiny billiard ball.
Thomson's model (1897). J. J. Thomson studied cathode rays, glowing beams inside a glass tube with most of the air removed. The rays bent toward a positively charged plate, so they had to be made of negative particles. They were the same no matter which metal he used, and each particle was far lighter than the lightest atom (today we know an electron has about 1/1836 the mass of a hydrogen atom). These particles, the electrons, had to be parts of all atoms, so atoms were not indivisible after all. Since atoms are neutral overall, Thomson went on to picture (by 1904) electrons stuck in a ball of positive charge, like raisins in a bun or plums in a pudding.
The gold foil experiment
In 1909, working in Ernest Rutherford's lab in Manchester, Hans Geiger and Ernest Marsden fired fast, positively charged alpha particles from a radioactive source at a very thin sheet of gold foil. A screen around the foil flashed wherever an alpha particle hit. If Thomson's model were right, the spread-out positive charge would be too weak to turn the heavy, fast alpha particles, and they should all pass almost straight through. Most did. But a very few, roughly 1 in 8000, bounced off at large angles, some straight back. Rutherford said it was as if you fired a large artillery shell at tissue paper and it came back and hit you. In 1911 he published his explanation.
- Most alpha particles passed straight through the foil.
- Some were deflected (turned) slightly.
- A very few bounced back toward the source.
- An atom is mostly empty space.
- The positive charge is concentrated in a tiny region that repels positive alpha particles passing close to it.
- Almost all the atom's mass is packed into that tiny, dense, positive centre: the nucleus.
Bohr's model (1913). Rutherford's model had a problem: why don't the negative electrons spiral into the positive nucleus? And when an element such as hydrogen is heated, it gives off only certain colours of light, seen as separate bright lines (a line spectrum) rather than a full rainbow. The Danish physicist Niels Bohr explained both. In his model, electrons can only travel in certain fixed orbits, called energy levels or shells. An electron can jump up to a higher shell when it absorbs energy, and it gives off light of one exact colour when it drops back down. Because the shells have fixed energies, only certain colours appear.
Chadwick's neutron (1932). Helium has 2 protons but about 4 times the mass of hydrogen, which has 1 proton. Something else in the nucleus had to carry mass. The English physicist James Chadwick, a former student of Rutherford, bombarded beryllium with alpha particles and found a radiation that was not bent by electric or magnetic fields, so it had no charge, yet it could knock protons out of paraffin wax, so its particles had about the same mass as a proton. He called them neutrons.