Atomic Models: Thomson, Rutherford, Bohr Explained

Scientists once thought the atom was a solid, indivisible ball. Three models in a row — Thomson's, Rutherford's and Bohr's — each improved on the last by explaining new experimental evidence. Together they show how science works: a model survives only until an experiment it cannot explain comes along.

These models matter because every later idea about electrons, from orbitals to the periodic table, grew out of them. Exams love to ask which experiment killed which model, so knowing the evidence behind each one is the fastest way to full marks.

Key Concepts

• Thomson model (1904): the "plum pudding" model — electrons scattered through a sphere of uniform positive charge, like seeds in a pudding. It explained why atoms are electrically neutral and why cathode rays are the same for all elements.
• Rutherford model (1911): the nuclear model — almost all the atom's mass and all its positive charge sit in a tiny central nucleus, with electrons moving in the empty space around it.
• Bohr model (1913): electrons circle the nucleus only in certain allowed orbits with fixed energies; an electron absorbs or emits light only when it jumps between these orbits.
• Gold foil experiment (Geiger–Marsden, 1909): alpha particles were fired at thin gold foil. Most passed straight through, but about 1 in 8000 bounced back — impossible for a diffuse pudding of charge, but exactly what a small dense nucleus would do.
• Atomic spectra: each element emits light only at specific wavelengths, which the Thomson and Rutherford models could not explain but the Bohr model could.

Formulas

• Bohr's orbit condition: mvr = n·h/2π, where m is the electron mass, v its speed, r the orbit radius, h Planck's constant, and n = 1, 2, 3, … (the principal quantum number). Angular momentum is quantised in units of h/2π.
• Planck's constant: h = 6.626 × 10⁻³⁴ J·s.

Units: mvr has units of angular momentum, kg·m²/s (equivalently J·s). The formula applies only to hydrogen-like atoms with a single electron.

Worked Example

Calculate the angular momentum of an electron in the second Bohr orbit (n = 2) of hydrogen.

1. Use mvr = n·h/2π with n = 2.
2. Substitute h = 6.626 × 10⁻³⁴ J·s: L = 2 × 6.626 × 10⁻³⁴ / (2π).
3. Simplify: L = 6.626 × 10⁻³⁴ / π.
4. Compute: 6.626 / 3.1416 = 2.109, so L = 2.109 × 10⁻³⁴ J·s.

The angular momentum is 2.109 × 10⁻³⁴ J·s, exactly twice the value for n = 1.

Common Misconceptions

• The Bohr model is the final, correct picture of the atom. It works beautifully for hydrogen but fails for multi-electron atoms and cannot explain fine spectral lines; the quantum mechanical orbital model replaced it.
• Rutherford said electrons orbit like planets and that this was stable. Rutherford proposed the nucleus but offered no mechanism for electron motion; classical physics actually predicts orbiting electrons would radiate energy and spiral into the nucleus — the very failure Bohr fixed.
• Thomson's model was useless. It correctly established that electrons are universal parts of all atoms and that positive charge balances them; it was only replaced when the gold foil experiment demanded a concentrated nucleus.

Real-World Applications

• Atomic emission spectra from Bohr's idea are used to identify elements in stars and in lab samples (flame tests, spectroscopy).
• Rutherford's scattering principle is used in Rutherford backscattering spectrometry to analyse thin films in the semiconductor industry.
• Understanding atomic structure underpins X-ray imaging, lasers and all of modern electronics.
• The quantised-orbit idea led directly to quantum numbers and the periodic table's layout.

Thomson spread the charge through the atom, Rutherford packed it into a nucleus after the gold foil results, and Bohr froze electrons into fixed orbits to explain spectral lines. Each model fixed the previous one's failure — and each failure pointed the way to modern quantum theory.