Electron subshell
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Electron_subshell".

Example of a sodium electron shell model

An electron shell may be crudely thought of as an orbit followed by electrons around an atom nucleus. Because each shell can contain only a fixed number of electrons, each shell is associated with a particular range of electron energy, and thus each shell must fill completely before electrons can be added to an outer shell. The electrons in the outermost shell determine the chemical properties of the atom (see Valence shell). For an explanation of why electrons exist in these shells see electron configuration.¹

History

The existence of electron shells was first observed experimentally in Charles Barkla's and Henry Moseley's X-ray absorption studies. Barkla labeled them with the letters K, L, M, N, O, P, and Q. The origin of this terminology was alphabetic. A J series was also suspected, though later experiments indicated that the K absorption lines are produced by the innermost electrons. 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 is derived 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.

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 determining 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 one or more subshells, which are themselves composed of atomic orbitals. For example, the first (K) shell has one subshell, called "1s"; the second (L) shell has two subshells, called "2s" and "2p"; the third shell has "3s", "3p", and "3d"; and so on.¹ The various possible subshells are shown in the following table:

• The first column is the "subshell label", a lowercase-letter label for the type of subshell. For example, the "4s subshell" is a subshell of the fourth (N) shell, with the type ("s") described in the first row.
• The second column is the azimuthal quantum number of the subshell. The precise definition involves quantum mechanics, but it is a number that characterizes the subshell.
• The third column is the maximum number of electrons that can be put into a subshell of that type. For example, the top row says that each s-type subshell ("1s", "2s", etc.) can have at most two electrons in it.
• The fourth column says which shells have a subshell of that type. For example, looking at the top two rows, every shell has an s subshell, while only the second shell and higher have a p subshell (i.e., there is no "1p" subshell).
• The final column gives the historical origin of the labels s, p, d, and f. They come from early studies of atomic spectral lines.

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 exactly the same level of energy,² with later subshells having more energy per electron than earlier ones. This effect is great enough that the energy ranges associated with shells can overlap (see Valence shells and Aufbau Principle).

Number of electrons in each shell

• Each s subshell holds at most 2 electrons

• Each p subshell holds at most 6 electrons

• Each d subshell holds at most 10 electrons

• Each f subshell holds at most 14 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.³⁴ 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

Main article: Valence electron

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. While atoms with complete valence shells (noble gases) are the most chemically non-reactive, 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.⁵

However, the 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 that the 3d electrons have more energy than 4s electrons, and are therefore more important in chemical reactions, hence making them valence electrons although they are not in the so-called valence shell.⁶

See also

• Electron counting
• 18-Electron rule

References

Wikimedia Commons has media related to: Electron shell diagrams