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Our modern picture of the atom — a tiny, dense nucleus surrounded by electrons in shells — was not discovered in one flash of insight. It was built up over more than a century, as scientists gathered new evidence and altered the model to fit. That story is a perfect example of how scientific models develop: a model is accepted while it explains the evidence, then revised or replaced when new experiments reveal something it cannot account for. This lesson, part of Topic C1 of your OCR Gateway Combined Science course, traces the model from Dalton's solid spheres to the nuclear atom, and shows how the famous alpha-scattering experiment led to the discovery of the nucleus.
By the end of this lesson you should be able to describe each stage in the development of the atomic model, name the scientist and the evidence behind each change, explain how the alpha-scattering experiment led to the nuclear model, and explain why scientific models change over time.
This lesson develops AO1 (recalling each stage of the atomic model and the scientists behind it) and AO3 (analysing how the alpha-scattering evidence forced the plum-pudding model to be replaced by the nuclear model, and evaluating why scientific models change over time).
flowchart LR
A["Dalton<br/>early 1800s<br/>solid spheres"] --> B["Thomson<br/>1897<br/>plum-pudding model"]
B --> C["Rutherford<br/>1909-1911<br/>nuclear model"]
C --> D["Bohr<br/>1913<br/>electron shells"]
D --> E["Chadwick<br/>1932<br/>neutron discovered"]
Each step kept what still worked from the one before and changed only what the new evidence demanded. We will take them in turn.
John Dalton proposed that all matter is made of tiny solid spheres that cannot be divided or broken — he called them atoms, from a Greek word meaning "uncuttable". In his model:
Dalton's atoms were featureless solid balls with no internal structure — no protons, neutrons or electrons, because none of these had yet been discovered. The model explained a great deal about how elements combine — for instance, why elements always react together in fixed, whole-number ratios, which makes perfect sense if matter comes in indivisible units — and it was a huge step forward, but it was incomplete.
It is important to see that Dalton's model was not so much "wrong" as as far as the evidence of the time could take it. In the early 1800s no experiment had yet shown that atoms contained anything smaller, so picturing them as solid, indivisible spheres was a perfectly sensible reading of what was known. This is the pattern you will see repeated through the lesson: each model was the best available explanation of the evidence at the time, and each was revised only when new evidence appeared that the old picture could not handle. Judging earlier scientists by what we know now misses the point — they were reasoning correctly from the data they had.
Exam Tip: Remember Dalton for the idea of atoms as tiny, solid, indivisible spheres with no internal parts. The later discovery of charged particles is exactly what showed this was too simple, so the model had to be revised.
In 1897, J. J. Thomson discovered the electron — a tiny, negatively charged particle far smaller than an atom. This was the first evidence that atoms are not solid, indivisible spheres but have smaller parts inside them.
Because atoms are neutral overall, Thomson reasoned that the negative electrons must be balanced by some positive charge. He proposed the plum-pudding model: the atom is a ball of positive charge with the tiny negative electrons embedded in it, like currants (plums) dotted through a pudding. There was no nucleus in this model — the positive charge was spread out through the whole atom.
Exam Tip: Two facts about Thomson are commonly tested: he discovered the electron (1897), and his plum-pudding model had electrons embedded in a ball of positive charge, with no nucleus.
The plum-pudding model was overturned by an experiment carried out by Ernest Rutherford with his colleagues Hans Geiger and Ernest Marsden (around 1909), often called the alpha-scattering or gold-foil experiment.
A beam of alpha particles (small, positively charged particles) was fired at a very thin sheet of gold foil, only a few atoms thick. A detector around the foil recorded where the alpha particles went after passing through — or bouncing off — it.
The thinking behind the experiment is worth understanding, because it is the key to the whole story. Alpha particles are positively charged and move very fast. If the plum-pudding model were correct, the positive charge of each gold atom would be spread thinly through the whole atom, far too diffuse to deflect a fast, heavy alpha particle by much. The prediction was therefore clear: the alpha particles should all pass almost straight through the foil, with at most very small deflections, as if passing through a thin fog. The experiment was designed to test exactly this prediction — and the fact that it failed is what made the result so important.
In fact, three things happened:
| Observation | Conclusion |
|---|---|
| Most alpha particles passed straight through the foil undeflected | Most of the atom is empty space |
| A small number were deflected through large angles | The atom contains a concentrated positive charge that repels the positive alpha particles |
| A very few (about 1 in 8000) bounced almost straight back | The positive charge (and the mass) is concentrated in a tiny, dense region — the nucleus |
These results could not be explained by the plum-pudding model. Rutherford concluded that the atom must have a tiny, dense, positively charged nucleus at its centre, holding most of the atom's mass, with the electrons surrounding it and most of the atom being empty space. This is the nuclear model of the atom. A little later, Rutherford also showed that the nucleus contains positively charged particles, which were named protons.
Exam Tip: The alpha-scattering observation → conclusion mapping is a classic exam question. Learn the three pairs: most pass through → mostly empty space; some deflected → concentrated positive charge; a few bounce back → tiny, dense nucleus holding most of the mass.
Rutherford's nuclear model had a problem: it did not explain why the negative electrons did not simply spiral inwards and crash into the positive nucleus. In 1913, Niels Bohr refined the model. He proposed that the electrons orbit the nucleus at fixed distances, in specific energy levels (or shells), rather than anywhere around it. Electrons in these fixed shells are stable and do not spiral in. Bohr's idea of electrons in shells matches experimental evidence and is the model you use for electronic structure today.
Exam Tip: Bohr's contribution was the idea that electrons occupy fixed energy levels (shells) at set distances from the nucleus — this is the shell model used for working out electronic structures.
The last major piece was the neutron. The known protons did not account for all of the mass of the nucleus — the nucleus was heavier than its protons alone could explain — which suggested that another, uncharged particle was also present. In 1932, James Chadwick provided experimental evidence for the neutron — a particle with a similar mass to the proton but no charge, found in the nucleus. This completed the modern picture: a nucleus of protons and neutrons surrounded by electrons in shells.
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