The Bohr model of the atom

Rutherford (1911) pictured the atom as a positive nucleus with electrons circling it, but this model faced two problems: by classical electrodynamics an accelerating, orbiting electron should radiate, lose energy and fall into the nucleus; and a continuous spectrum was predicted, while atoms in experiment give line spectra. In 1913 Bohr offered a solution with two postulates. First: an atom has stationary states in which the electron does not radiate; in the corresponding orbits angular momentum is quantised, m_e·υ·r = nħ, ħ = h/2π, n = 1, 2, 3, … Second: when an atom passes from one state to another it emits or absorbs one photon, hν = E_n − E_m. For hydrogen r_n = n²r_B (r_B = 0.529·10⁻¹⁰ m, the Bohr radius) and E_n = −13.6 eV/n². The energy is negative: the electron is bound; n = 1 is the ground state, n > 1 are excited states; ionising an atom from the ground state takes 13.6 eV. A jump from a higher level down emits a photon, the reverse absorbs one; hence the spectrum consists of lines at certain frequencies. Bohr's theory works well only for hydrogen-like atoms and was later replaced by quantum mechanics.