The theory that preceded Bohr's model was the Rutherford model of the atom, which postulated that an atom contained a core nucleus and a number of electrons orbiting around it. This model was somewhat limited, and couldn't explain the emission spectra observed from a hydrogen lamp.
Max Planck had previously put forward his theory of quantum physics, in which energy could only be absorbed and emitted in discrete packets or quanta. Bohr applied this theory to his model of the atom, postulating that electrons could only exist in certain fixed energy orbits around the nucleus, and that they could absorb and emit specific amounts of energy to progress between.
The Bohr model contains a center nucleus containing protons and neutrons. Electrons orbit the nucleus in specific orbits with a fixed radius, with each orbit having a principle atomic number of n = 1, 2, 3 with successive numbers indicating an orbital with a higher radius and electron energy, and with the energy gap between the first and second being larger than the second and third and so on. Electrons would initially fill the lower energy orbits before the higher ones, this formation being known as the ground state.
If a electron in the ground state absorbed energy it would be able to be promoted to a higher energy orbit. After time it would fall back to its original state releasing a photon of light with energy corresponding to the gap between the two levels. The electron did not necessarily fall straight back, for example it could jump from n=1 to n=3 and either fall back from n=3 to n=1 or from n=3 to n=2 and then n=2 to n=1. As the energy levels between each level were different this allowed for a number of wavelengths to be emitted if a sample was exposed to a large energy source, however as the values were fixed only specific wavelengths were allowed hence explaining the discrete lines on the emission spectra of elements.
Bohr modeled his atom on the emission spectrum of hydrogen. The elemental atom of hydrogen at ground state contains an electron in the first electron orbit, n=1. When the atom is excited the electron can jump into a higher state including energy levels n=2, n=3, n=4, n=5, n=6, and then would fall back to its ground state. Bohr calculated a number of transitions were possible, including:
- Paschen Series - wavelength in the IR range, when an electron falls to n=3 from a higher orbit
- Balmer Series - wavelength in the visible range, when an electron falls to n=2 from a higher orbit, specifically
- n=3 to n=2 - 656.3nm (red)
- n=4 to n=2 - 486.1nm (blue)
- n=5 to n=2 - 434.1nm (violet)
- n=6 to n=2 - 410.2nm (violet)
- n=7 to n=2 - 397.0nm (violet)
- n=8 to n=2 - 388.9nm (violet)
- Lymann Series - wavelength in the UV range, when an electron falls from n=2 to n=1
Although providing a revolutionary approach to scientific understanding of the atom the Bohr model was limited in several aspects:
- Although it allowed calculations of the energy orbits around hydrogen these couldn't be applied to other atoms, it was later found this was because different atoms have different electron energy levels
- It could not explain the concept of doublets, these were shown to be due to subshells within energy levels
- It contained electrons orbiting with fixed radii, this was later shown to be false by the Heisenberg uncertainty principle, an effect of which is that the probability of an electron appearing in a location can only be predicted, its path cannot be mapped