Periodic table with electron shells

An electron shell may be thought of as an orbit followed by electrons around an atom's nucleus. The closest shell to the nucleus is called the "1 shell" (also called "K shell"), followed by the "2 shell" (or "L shell"), then the "3 shell" (or "M shell"), and so on further and further from the nucleus. The shell letters K, L, M, ... are alphabetical.

Each shell can contain only a fixed number of electrons: The 1st shell can hold up to two electrons, the 2nd shell can hold up to eight electrons, the 3rd shell can hold up to 18, and so on. Since electrons are electrically attracted to the nucleus, an atom's electrons will generally occupy outer shells only if the more inner shells have already been completely filled by other electrons. However, this is not a strict requirement: Atoms may have two or even three outer shells that are only partly filled with electrons. (See Madelung rule for more details.) For an explanation of why electrons exist in these shells see electron configuration.^[1]

The electrons in the partially filled outermost shell (or shells) determine the chemical properties of the atom; it is called valence shell.

Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals.

History[link]

The shell terminology comes from Arnold Sommerfeld's modification of the Bohr model. Sommerfeld retained Bohr's planetary model, but added mildly elliptical orbits (characterized by additional quantum numbers l and m) to explain the fine spectroscopic structure of some elements.^[2] The multiple electrons with the same principal quantum number (n) had close orbits that formed a "shell" of finite thickness instead of the infinitely thin circular orbit of Bohr's model.

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 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[link]

The electron shells are labeled 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 farther from the nucleus than those in inner shells. This makes them more important in determining how the atom reacts chemically and behaves as a conductor, because the pull of the atom's nucleus upon them is weaker and more easily broken. In this way, a given element's reactivity is highly dependent upon its electronic configuration.

Subshells[link]

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. In each case the figure is 4 greater than the one above 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. The other labels, namely g, h and i, are an alphabetic continuation following the last historically originated label of f.

Although it is commonly stated that all the electrons in a shell have the same energy, this is an approximation. However, the electrons in one 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[link]

An atom's electron shells are filled according to the following theoretical constraints:

• 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
• Each g subshell holds at most 18 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 is only achieved (by known elements) for the first four shells (K,L,M,N). 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[link]

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. Just as 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.^[7]

However, this is a simplification of the truth. The electrons that determine how an atom reacts chemically are those that travel farthest from the nucleus, that is, 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, making them valence electrons although they are not in the so-called valence shell.^[8]

List of elements with electrons per shell[link]

The list below gives the elements arranged by increasing atomic number and shows the number of electrons per shell. At a glance, one can see that subsets of the list show obvious patterns. In particular, the seven elements (in light blue) before a noble gas (group 18, in yellow) higher than helium have the number of electrons in the valence shell in arithmetic progression. (However, this pattern may break down in the seventh period due to relativistic effects.)

Sorting the table by chemical group shows additional patterns, especially with respect to the last two outermost shells. (Elements 57 to 71 belong to the lanthanides, while 89 to 103 are the actinides.)

The list below is primarily consistent with the Aufbau rule. However, there are a number of exceptions to the rule; for example palladium (atomic number 46) has no electrons in the fifth shell, unlike other atoms with lower atomic number. Some entries in the table are uncertain, when experimental data is unavailable. (For example, some atoms have such short half-life that it is impossible to measure their electron configurations).

See also[link]

• Electron counting
• 18-Electron rule

References[link]

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