لایه الکترونی (به انگلیسی: Electron Shell) یا مدار الکترونی در علم شیمی اصطلاحی برای تعریف مدار یا مدارهایی در اطراف هستهٔاتم است که الکترونها در حوزهٔ آن، در اطراف هستهٔ اتم، در نظر گرفته میشوند. نزدیکترین لایه به هسته، لایه نخست، دومین لایه، لایه دوم و به همین صورت تا لایه هفتم،میتوان برای تعداد لایهها متصور بود.
هر لایه، به تعداد ثابتی میتواند الکترون را در خود جای دهد. لایه نخست توانایی جای دادن حداکثر دو الکترون و لایه دوم تا حداکثر هشت الکترون را در خود جای میدهد. تعداد الکترونها در لایههای بعدی، از طریق رابطه ۲n۲ معین میشوند. در این رابطه، n نمایانگر عدد لایه الکترونی است.
In chemistry and atomic physics, an electron shell, or a principal energy level[clarification needed], 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 farther and farther from the nucleus. The shells correspond with the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).
Each shell can contain only a fixed number of electrons: The first shell can hold up to two electrons, the second shell can hold up to eight (2 + 6) electrons, the third shell can hold up to 18 (2 + 6 + 10) and so on. The general formula is that the nth shell can in principle hold up to 2(n2) electrons. 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 incomplete outer shells. (See Madelung rule for more details.) For an explanation of why electrons exist in these shells see electron configuration.
The electrons in the outermost occupied shell (or shells) determine the chemical properties of the atom; it is called the valence shell.
Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals.
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 numbersℓ and m) to explain the fine spectroscopic structure of some elements. The multiple electrons with the same principal quantum number (n) had close orbits that formed a "shell" of positive 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 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.
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; the fourth shell has 4s, 4p, 4d and 4f; the fifth shell has 5s, 5p, 5d, and 5f and can theoretically hold more in the 5g subshell that is not occupied in the ground-state electron configuration of any known element. 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 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.
For multielectron atoms n is a poor indicator of electron's energy. Energy spectra of some shells interleave.
Subshell max electrons
Shell max electrons
2 + 6 = 8
2 + 6 + 10 = 18
2 + 6 + 10 + 14 = 32
2 + 6 + 10 + 14 + 18 = 50
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, 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 subshell is constrained to hold 4ℓ+ 2 electrons at most, namely:
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; in general, the nth shell can 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. 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).
The valence shell is the outermost shell of an atom. Valence electrons in s- and p-elements reside in this shell. Such elements with complete valence shells (noble gases) are the least chemically reactive, while those with only one electron in their valence shells (alkali metals, with the exception of hydrogen) or just missing one electron from having a complete shell (halogens) are the most reactive. Hydrogen, being the only element which has one electron in its valence shell but is also just missing one electron from having a complete shell, has unique properties.
However, this terminology is somewhat misleading in the case of transition and inner transition metals. In these elements, a valence electron can also be in an inner shell. Thus, the electrons that determine how an atom reacts chemically are those that travel farthest from the nucleus, that is, those with the highest energy, and not necessarily in the valence shell.
List of elements with electrons per shell
The list below gives the elements arranged by increasing atomic number and shows the number of electrons per shell. At a glance, the subsets of the list show obvious patterns. In particular, every set of seven elements (in electric blue) before each noble gas (group 18, in yellow) heavier than helium have the number of electrons in the valence shell in arithmetic progression, namely one to seven. (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 principle. 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, the elements past 108 have such short half-lives that their electron configurations have not yet been measured.)
^The statement that the electrons in one 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.
^Orbitals. Chem4Kids. Retrieved on 1 December 2011.