Editing Atomic orbital (section)

=== Bohr atom ===
In 1909, [[Ernest Rutherford]] discovered that the bulk of the atomic mass was tightly condensed into a nucleus, which was also found to be positively charged. It became clear from his analysis in 1911 that the plum pudding model could not explain atomic structure. In 1913, Rutherford's post-doctoral student, [[Niels Bohr]], proposed a new model of the atom, wherein electrons orbited the nucleus with classical periods, but were permitted to have only discrete values of angular momentum, quantized in units [[Planck constant|ħ]].<ref name="Bohr 1913 476" /> This constraint automatically allowed only certain electron energies. The [[Bohr model]] of the atom fixed the problem of energy loss from radiation from a ground state (by declaring that there was no state below this), and more importantly explained the origin of spectral lines.

[[File:Bohr atom model.svg|thumb|The [[Bohr model|Rutherford–Bohr model]] of the hydrogen atom]]
After Bohr's use of [[Albert Einstein|Einstein]]'s explanation of the [[photoelectric effect]] to relate energy levels in atoms with the wavelength of emitted light, the connection between the structure of electrons in atoms and the [[Emission spectra|emission]] and [[absorption spectra]] of atoms became an increasingly useful tool in the understanding of electrons in atoms. The most prominent feature of emission and absorption spectra (known experimentally since the middle of the 19th century), was that these atomic spectra contained discrete lines. The significance of the Bohr model was that it related the lines in emission and absorption spectra to the energy differences between the orbits that electrons could take around an atom. This was, however, ''not'' achieved by Bohr through giving the electrons some kind of wave-like properties, since the idea that electrons could behave as [[matter waves]] was not suggested until eleven years later. Still, the Bohr model's use of quantized angular momenta and therefore quantized energy levels was a significant step toward the understanding of electrons in atoms, and also a significant step towards the development of [[quantum mechanics]] in suggesting that quantized restraints must account for all discontinuous energy levels and spectra in atoms.

With [[Louis de Broglie|de Broglie]]'s suggestion of the existence of electron matter waves in 1924, and for a short time before the full 1926 [[Schrödinger equation]] treatment of [[hydrogen-like atom]]s, a Bohr electron "wavelength" could be seen to be a function of its momentum; so a Bohr orbiting electron was seen to orbit in a circle at a multiple of its half-wavelength. The Bohr model for a short time could be seen as a classical model with an additional constraint provided by the 'wavelength' argument. However, this period was immediately superseded by the full three-dimensional wave mechanics of 1926. In our current understanding of physics, the Bohr model is called a semi-classical model because of its quantization of angular momentum, not primarily because of its relationship with electron wavelength, which appeared in hindsight a dozen years after the Bohr model was proposed.

The Bohr model was able to explain the emission and absorption spectra of [[hydrogen]]. The energies of electrons in the ''n''&nbsp;=&nbsp;1, 2, 3, etc. states in the Bohr model match those of current physics. However, this did not explain similarities between different atoms, as expressed by the periodic table, such as the fact that [[helium]] (two electrons), neon (10 electrons), and [[argon]] (18 electrons) exhibit similar chemical inertness. Modern [[quantum mechanics]] explains this in terms of [[electron shell]]s and subshells which can each hold a number of electrons determined by the [[Pauli exclusion principle]]. Thus the ''n''&nbsp;=&nbsp;1 state can hold one or two electrons, while the ''n'' = 2 state can hold up to eight electrons in 2s and 2p subshells. In helium, all ''n''&nbsp;=&nbsp;1 states are fully occupied; the same is true for ''n''&nbsp;=&nbsp;1 and ''n''&nbsp;=&nbsp;2 in neon. In argon, the 3s and 3p subshells are similarly fully occupied by eight electrons; quantum mechanics also allows a 3d subshell but this is at higher energy than the 3s and 3p in argon (contrary to the situation for hydrogen) and remains empty.