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Orbital motions of electrons
Although it is common in diagrams to show electrons as objects following exact orbits like miniature planets, they actually move randomly and their so-called "orbits" represent only an average position. [2]
Shells
The electron shells are labelled K, L, M, N, O, P, and Q; or 1, 2, 3, 4, 5, 6, and 7; going from innermost shell outwards. Electrons in outer shells have higher average energy and travel further from the nucleus than those in inner shells, making them more important in determing how the atom reacts chemically and behaves as a conductor, etc, because the pull of the atom's nucleus upon them is weaker and more easily broken.
Subshells
Each shell is composed of subshells labeled s,p,d,and f; which are themselves composed of orbitals. An s subshell orbital can contain a maximum of two electrons. The first principle shell, K (or 1), has one subshell, the 1s shell. The second principal shell has two subshells, 2s and 2p. The third shell subshells designated 3s, 3p, and 3d, etc. [3]
Although it is commonly stated that all the electrons in a shell have the same energy, this is an approximation. However, the electrons in a subshell do have the same level of energy, with later subshells having more energy per electron that earlier ones. This effect is great enough that the energy ranges associated with shells can overlap - see "Valence shells" below.
Number of electrons in each shell
- Each s subshell holds no more than two electrons
- Each p subshell holds no more than six electrons
- Each d subshell holds no more than ten electrons
- Each f subshell holds no more than fourteen electrons
Therefore, the K shell, which contains only an s subshell, can hold up to 2 electrons; the L shell, which contains an s and a p, can hold up to 2+6=8 electrons; and so forth. The general formula is that the nth shell can in principle hold up to 2n2 electrons.
Although that formula gives the maximum in principle, in fact that maximum can only be achieved (by known elements) for the first four shells (K,L,M,N). In fact, no known element has more than 32 electrons in any one shell.[4][5] This is because the subshells are filled according to the Aufbau principle. The first elements to have more than 32 electrons in one shell would belong to the g-block of period 8 of the periodic table. These elements would have some electrons in their 5g subshell and thus have more than 32 electrons in the O shell (fifth principal shell).
Valence shells
The valence shell is the outermost shell of an atom. It is usually (and misleadingly) said that the electrons in this shell make up its valence electrons, that is, the electrons that determine how the atom behaves in chemical reactions, where atoms with complete valence shells are the most chemical non-reactive, while those with only one electron in their valence shells (alkalis) or just missing one electron from having a complete shell (halogens) are the most reactive. [6]
However, the real truth is more complicated. The electrons that determine how an atom reacts chemically are those that travel furthest from the nucleus - i.e. those with the most energy. As stated in "Subshells", electrons in the inner subshells have less energy than those in outer subshells. This effect is great enough so that the 3d electrons have more energy than 4s electrons, and are hence more important in chemical reactions, hence making them valence electrons although they are not in the so-called valence shell. [7]
See also Electron counting and 18-Electron rule
History
The existence of electron shells was first observed experimentally in Charles Barkla's and Henry Moseley's X-ray absorption studies. Barkla labelled them with the letters K, L, M, N, O, P, and Q. (The origin of this terminology was alphabetic. K for hypothetical spectral lines that were never discovered.) These letters were later found to correspond to the n-values 1, 2, 3, etc. They are used in the spectroscopic Siegbahn notation.
The name for electron shells originates from the Bohr model, in which groups of electrons were believed to orbit the nucleus at certain distances, so that their orbits formed "shells" around the nucleus.
The physical chemist Gilbert Lewis was responsible for much of the early development of the theory of the participation of valence shell electrons in chemical bonding. Linus Pauling later generalized and extended the theory while applying insights from quantum mechanics.
References
