Electron shell: Difference between revisions
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[[Image:Electron shell 011 Sodium.svg|thumb|Example of a sodium electron shell model]] |
[[Image:Electron shell 011 Sodium.svg|thumb|Example of a sodium electron shell model]] |
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An '''electron shell''' may be crudely thought of as an [[orbit]] followed by [[electrons]] around an [[atom]] [[Atomic nucleus|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 |
An '''electron shell''' may be crudely thought of as Andrew's Poop an [[orbit]] followed by [[electrons]] around an [[atom]] [[Atomic nucleus|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 |
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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_electron|Valence shell]]). For an explanation of why electrons exist in these shells see [[electron configuration]].<ref>[http://www.corrosionsource.com/handbook/periodic/e_subshells.htm Corrosion Source - Electron Subshells]</ref> |
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_electron|Valence shell]]). For an explanation of why electrons exist in these shells see [[electron configuration]].<ref>[http://www.corrosionsource.com/handbook/periodic/e_subshells.htm Corrosion Source - Electron Subshells]</ref> |
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Revision as of 15:55, 15 September 2008
![](http://upload.wikimedia.org/wikipedia/commons/thumb/7/76/Electron_shell_011_Sodium.svg/220px-Electron_shell_011_Sodium.svg.png)
An electron shell may be crudely thought of as Andrew's Poop 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 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.
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 height.[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 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.[3] The various possible subshells are shown in the following table:
Subshell label | l | Max electrons | Shells containing it | Historical name |
---|---|---|---|---|
s | 0 | 2 | Every shell | sharp |
p | 1 | 6 | 2nd shell and higher | principal |
d | 2 | 10 | 3rd shell and higher | diffuse |
f | 3 | 14 | 4th shell and higher | fundamental |
g | 4 | 18 | 5th shell and higher | |
h | 5 | 22 | 6th shell and higher | |
i | 6 | 26 | 7th shell and higher |
- 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 it 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,[4] 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 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.[5][6] 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 chemically 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.[7]
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.[8]
See also
References
![](http://upload.wikimedia.org/wikipedia/en/thumb/4/4a/Commons-logo.svg/30px-Commons-logo.svg.png)
- ^ Corrosion Source - Electron Subshells
- ^ BBC - Electron Shells and Orbitals
- ^ Corrosion Source - Electron Subshells
- ^ The statement that the electrons in a subshell have exactly the same level of energy is true in an isolated atom, where it follows quantum-mechanically from the spherical symmetry of the system. When the atom is part of a molecule, this no longer holds; see, for example, crystal field theory.
- ^ Chem4Kids - Orbitals
- ^ Electron & Shell Configuration
- ^ Vision Learning - Chemical Reactions
- ^ Chemguide - Atomic Orbitals