The Bohr Model

Bohr's model was the first atomic model to integrate insights from quantum physics. The model was the first to explain the atomic spectrum of hydrogen, and it was widely accepted thanks to this success.

In Bohr's model, electrons orbit the nucleus in circular orbits, much like in the previous Rutherford atomic model. However, the main novelty introduced by Bohr is that electrons can't occupy orbits of just any size, only certain orbits of discrete sizes. The smaller the electron's orbit, the lower its energy level.

Electrons occupy these orbits and go around the nucleus on them without emitting any energy or radiation, but they emit or absorb energy (in the shape of a photon) when they move to a different orbit. The energy absorbed or emitted is the difference between the energy of the final orbit and that of the initial orbit:

In this equation, E_i and E_f are the initial and final orbital energies, respectively.

Because the energy is emitted as a photon (a light particle) it can be expressed as a function of the photon's frequency, that is, as Planck's constant h multiplied by the frequency of the photon v, or, equivalently, by the speed of light divided by the photon's wavelength.

Despite its success, Bohr's model had important limitations.

• It could not account for electron–electron interactions, and therefore it could not explain the structure of atoms with more than one electron, that is, any atom other than hydrogen.

• It could not explain why most energy levels in the atom have sub-levels.

• It contradicted the uncertainty principle, one of the core principles of quantum mechanics, which states that particles can't have well-defined positions and well-defined velocities at the same time, meaning that it is impossible to assign fixed paths to electrons like the orbits in Bohr's model.