Bohr Model

In Bohr's model, electrons don't spiral into the nucleus — they occupy discrete, stable orbits. Each orbit corresponds to a specific angular momentum: L = nℏ. The electron cannot exist between these shells. This was the first quantisation postulate in atomic physics.

The radius of each orbit grows quadratically: r_n = a₀ · n². The innermost shell (n = 1) sits at the Bohr radius a₀ ≈ 0.529 Å — half an angstrom from the nucleus.

When an electron jumps from a higher shell (n_i) to a lower one (n_f), the atom emits a photon. The photon's energy equals the difference between the two levels: ΔE = E_f − E_i.

This explains the line spectra of hydrogen — discrete wavelengths rather than a continuous rainbow. The Lyman, Balmer, and Paschen series correspond to transitions ending at n = 1, 2, and 3 respectively.

Bohr's model predicted the hydrogen spectrum with extraordinary accuracy — but it fails for multi-electron atoms. It treats electrons as classical particles on fixed orbits, lacking the probabilistic nature of quantum mechanics.

Schrödinger's wave equation replaced Bohr's orbits with probability clouds — orbitals where the electron is most likely to be found. Yet the Bohr model remains a powerful teaching tool and gives the correct energy levels for hydrogen.

→ Explore the quantum correction: Schrödinger Orbital Simulation