Preview
IB DP Physics — Unit E: Nuclear and Quantum Physics | SL & HL | E.1 Structure of the Atom
📋 What You Will Learn
- Observe how scattering experiments revealed the internal structure of the atom, moving from the Thomson model to the Rutherford nuclear model
- Define the key quantities describing atomic and nuclear structure: atomic number, mass number, nucleon number, and nuclear radius
- Describe the composition of the nucleus in terms of protons and neutrons, and explain the role of the strong nuclear force
- Distinguish between isotopes and explain how nuclear notation encodes atomic information
- Derive and apply the relationship between nuclear radius and nucleon number
- Analyse evidence from the Geiger–Marsden experiment and evaluate what it tells us about nuclear structure
🌍 Real-World Context Imagine firing a stream of tiny bullets at a sheet of tissue paper and finding that, occasionally, one bounces almost straight back at you. This is precisely what Geiger and Marsden observed in 1909 when they directed alpha particles at a thin gold foil — and it shattered everything physicists thought they knew about the atom. The result was so unexpected that Rutherford himself described it as "almost as incredible as if you fired 15-inch shells at a piece of tissue paper and they came back and hit you." That single experiment forced a complete reimagining of what matter is made of at its most fundamental level.
🔍 Guiding Question How did experimental evidence from alpha-particle scattering lead to the nuclear model of the atom, and what does the structure of the nucleus tell us about the forces and particles that hold matter together?
Physical Setup: The Geiger–Marsden Experiment
Before Rutherford's experiment, the dominant model of the atom was J.J. Thomson's "plum pudding" model, proposed in 1904. In this model, the atom was imagined as a diffuse, positively charged sphere — like a pudding — with negatively charged electrons embedded throughout it, like plums. The positive charge was spread uniformly across the entire atomic volume, so any particle passing through would experience only gentle, distributed deflections.
To test this model, Hans Geiger and Ernest Marsden, working under Rutherford at the University of Manchester, directed a beam of alpha particles (positively charged particles emitted by radioactive sources) at an extremely thin sheet of gold foil. They placed a zinc sulfide detector screen around the foil to observe where the alpha particles landed after passing through or being deflected by the foil.
The results were astonishing. The vast majority of alpha particles passed straight through the foil with little or no deflection — exactly as the plum pudding model predicted. However, a small fraction were deflected through large angles, and a tiny number (approximately 1 in 8000) bounced almost straight back. This was completely inconsistent with the plum pudding model, in which the positive charge was too diffuse to cause such strong deflections.
From Observation to Theory: The Nuclear Model
Rutherford reasoned that only a highly concentrated, massive, positively charged region at the centre of the atom could cause such large deflections. An alpha particle approaching this concentrated charge head-on would experience an enormous electrostatic repulsion, sufficient to reverse its direction entirely. The fact that most particles passed through undeflected told him that this central region must be extremely small compared to the overall size of the atom — most of the atom must be empty space.
🍑 Analogy — The Solar System Think of