Electron shell

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In chemistry and atomic physics, 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 the "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 to the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with the letters used in X-ray notation (K, L, M, …).

Contents

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. [1] For an explanation of why electrons exist in these shells see electron configuration. [2]

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

History

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 and m) to explain the fine spectroscopic structure of some elements. [3] 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.[ non-primary source needed ] Barkla labeled them with the letters K, L, M, N, O, P, and Q. [4] 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.

Subshells

3D views of some hydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown). Atomic-orbital-clouds spdf m0.png
3D views of some hydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown).

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. [2] The various possible subshells are shown in the following table:

Subshell label Max electronsShells containing itHistorical name
s02Every shell sharp
p162nd shell and higher principal
d2103rd shell and higher diffuse
f3144th shell and higher fundamental
g4185th shell and higher (theoretically)(next in alphabet after f, excluding i) [5]

Number of electrons in each shell

Each subshell is constrained to hold 4+ 2 electrons at most, namely:

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. [1]

Shell
name
Subshell
name
Subshell
max
electrons
Shell
max
electrons
K1s22
L2s22 + 6 = 8
2p6
M3s22 + 6 + 10
= 18
3p6
3d10
N4s22 + 6 +
10 + 14
= 32
4p6
4d10
4f14
O5s22 + 6 +
10 + 14 +
18 = 50
5p6
5d10
5f14
5g18

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. [6] [7] 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).

Subshell energies and filling order

For multielectron atoms n is a poor indicator of electron's energy. Energy spectra of some shells interleave. Atomic orbitals as triangles.svg
For multielectron atoms n is a poor indicator of electron's energy. Energy spectra of some shells interleave.
The states crossed by same red arrow have same
n
+
l
{\displaystyle n+\ell }
value. The direction of the red arrow indicates the order of state filling. Aufbau Principle.png
The states crossed by same red arrow have same value. The direction of the red arrow indicates the order of state filling.

Although it is sometimes 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.

The filling of the shells and subshells with electrons proceeds from subshells of lower energy to subshells of higher energy. This follows the n + ℓ rule which is also commonly known as the Madelung rule. Subshells with a lower n + ℓ value are filled before those with higher n + ℓ values. In the case of equal n + ℓ values, the subshell with a lower n value is filled first.

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 five elements (in   electric blue) before each noble gas (group 18, in   yellow) heavier than helium have successive numbers of electrons in the outermost shell, namely three to seven.

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.)

Z Element No. of electrons/shell Group
1 Hydrogen 11
2 Helium 218
3 Lithium 2, 11
4 Beryllium 2, 22
5 Boron 2, 313
6 Carbon 2, 414
7 Nitrogen 2, 515
8 Oxygen 2, 616
9 Fluorine 2, 717
10 Neon 2, 818
11 Sodium 2, 8, 11
12 Magnesium 2, 8, 22
13 Aluminium 2, 8, 313
14 Silicon 2, 8, 414
15 Phosphorus 2, 8, 515
16 Sulfur 2, 8, 616
17 Chlorine 2, 8, 717
18 Argon 2, 8, 818
19 Potassium 2, 8, 8, 11
20 Calcium 2, 8, 8, 22
21 Scandium 2, 8, 9, 23
22 Titanium 2, 8, 10, 24
23 Vanadium 2, 8, 11, 25
24 Chromium 2, 8, 13, 16
25 Manganese 2, 8, 13, 27
26 Iron 2, 8, 14, 28
27 Cobalt 2, 8, 15, 29
28 Nickel 2, 8, 16, 210
29 Copper 2, 8, 18, 111
30 Zinc 2, 8, 18, 212
31 Gallium 2, 8, 18, 313
32 Germanium 2, 8, 18, 414
33 Arsenic 2, 8, 18, 515
34 Selenium 2, 8, 18, 616
35 Bromine 2, 8, 18, 717
36 Krypton 2, 8, 18, 818
37 Rubidium 2, 8, 18, 8, 11
38 Strontium 2, 8, 18, 8, 22
39 Yttrium 2, 8, 18, 9, 23
40 Zirconium 2, 8, 18, 10, 24
41 Niobium 2, 8, 18, 12, 15
42 Molybdenum 2, 8, 18, 13, 16
43 Technetium 2, 8, 18, 13, 27
44 Ruthenium 2, 8, 18, 15, 18
45 Rhodium 2, 8, 18, 16, 19
46 Palladium 2, 8, 18, 1810
47 Silver 2, 8, 18, 18, 111
48 Cadmium 2, 8, 18, 18, 212
49 Indium 2, 8, 18, 18, 313
50 Tin 2, 8, 18, 18, 414
51 Antimony 2, 8, 18, 18, 515
52 Tellurium 2, 8, 18, 18, 616
53 Iodine 2, 8, 18, 18, 717
54 Xenon 2, 8, 18, 18, 818
55 Caesium 2, 8, 18, 18, 8, 11
56 Barium 2, 8, 18, 18, 8, 22
57 Lanthanum 2, 8, 18, 18, 9, 2
58 Cerium 2, 8, 18, 19, 9, 2
59 Praseodymium 2, 8, 18, 21, 8, 2
60 Neodymium 2, 8, 18, 22, 8, 2
61 Promethium 2, 8, 18, 23, 8, 2
62 Samarium 2, 8, 18, 24, 8, 2
63 Europium 2, 8, 18, 25, 8, 2
64 Gadolinium 2, 8, 18, 25, 9, 2
65 Terbium 2, 8, 18, 27, 8, 2
66 Dysprosium 2, 8, 18, 28, 8, 2
67 Holmium 2, 8, 18, 29, 8, 2
68 Erbium 2, 8, 18, 30, 8, 2
69 Thulium 2, 8, 18, 31, 8, 2
70 Ytterbium 2, 8, 18, 32, 8, 2
71 Lutetium 2, 8, 18, 32, 9, 23
72 Hafnium 2, 8, 18, 32, 10, 24
73 Tantalum 2, 8, 18, 32, 11, 25
74 Tungsten 2, 8, 18, 32, 12, 26
75 Rhenium 2, 8, 18, 32, 13, 27
76 Osmium 2, 8, 18, 32, 14, 28
77 Iridium 2, 8, 18, 32, 15, 29
78 Platinum 2, 8, 18, 32, 17, 110
79 Gold 2, 8, 18, 32, 18, 111
80 Mercury 2, 8, 18, 32, 18, 212
81 Thallium 2, 8, 18, 32, 18, 313
82 Lead 2, 8, 18, 32, 18, 414
83 Bismuth 2, 8, 18, 32, 18, 515
84 Polonium 2, 8, 18, 32, 18, 616
85 Astatine 2, 8, 18, 32, 18, 717
86 Radon 2, 8, 18, 32, 18, 818
87 Francium 2, 8, 18, 32, 18, 8, 11
88 Radium 2, 8, 18, 32, 18, 8, 22
89 Actinium 2, 8, 18, 32, 18, 9, 2
90 Thorium 2, 8, 18, 32, 18, 10, 2
91 Protactinium 2, 8, 18, 32, 20, 9, 2
92 Uranium 2, 8, 18, 32, 21, 9, 2
93 Neptunium 2, 8, 18, 32, 22, 9, 2
94 Plutonium 2, 8, 18, 32, 24, 8, 2
95 Americium 2, 8, 18, 32, 25, 8, 2
96 Curium 2, 8, 18, 32, 25, 9, 2
97 Berkelium 2, 8, 18, 32, 27, 8, 2
98 Californium 2, 8, 18, 32, 28, 8, 2
99 Einsteinium 2, 8, 18, 32, 29, 8, 2
100 Fermium 2, 8, 18, 32, 30, 8, 2
101 Mendelevium 2, 8, 18, 32, 31, 8, 2
102 Nobelium 2, 8, 18, 32, 32, 8, 2
103 Lawrencium 2, 8, 18, 32, 32, 8, 33
104 Rutherfordium 2, 8, 18, 32, 32, 10, 24
105 Dubnium 2, 8, 18, 32, 32, 11, 25
106 Seaborgium 2, 8, 18, 32, 32, 12, 26
107 Bohrium 2, 8, 18, 32, 32, 13, 27
108 Hassium 2, 8, 18, 32, 32, 14, 28
109 Meitnerium 2, 8, 18, 32, 32, 15, 2 (?)9
110 Darmstadtium 2, 8, 18, 32, 32, 16, 2 (?)10
111 Roentgenium 2, 8, 18, 32, 32, 17, 2 (?)11
112 Copernicium 2, 8, 18, 32, 32, 18, 2 (?)12
113 Nihonium 2, 8, 18, 32, 32, 18, 3 (?)13
114 Flerovium 2, 8, 18, 32, 32, 18, 4 (?)14
115 Moscovium 2, 8, 18, 32, 32, 18, 5 (?)15
116 Livermorium 2, 8, 18, 32, 32, 18, 6 (?)16
117 Tennessine 2, 8, 18, 32, 32, 18, 7 (?)17
118 Oganesson 2, 8, 18, 32, 32, 18, 8 (?)18

See also

Related Research Articles

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Atomic radius

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Quantum number Notation for conserved quantities in physics and chemistry

In chemistry and quantum physics, quantum numbers describe values of conserved quantities in the dynamics of a quantum system. Quantum numbers correspond to eigenvalues of operators that commute with the Hamiltonian—quantities that can be known with precision at the same time as the system's energy—and their corresponding eigenspaces. Together, a specification of all of the quantum numbers of a quantum system fully characterize a basis state of the system, and can in principle be measured together.

In quantum mechanics, the principal quantum number is one of four quantum numbers assigned to each electron in an atom to describe that electron's state. Its values are natural numbers making it a discrete variable.

Azimuthal quantum number

The azimuthal quantum number is a quantum number for an atomic orbital that determines its orbital angular momentum and describes the shape of the orbital. The azimuthal quantum number is the second of a set of quantum numbers which describe the unique quantum state of an electron. It is also known as the orbital angular momentum quantum number, orbital quantum number or second quantum number, and is symbolized as .

The magnetic quantum number is one of four quantum numbers in atomic physics. The set is: principal quantum number, azimuthal quantum number, magnetic quantum number, and spin quantum number. Together, they describe the unique quantum state of an electron. The magnetic quantum number distinguishes the orbitals available within a subshell, and is used to calculate the azimuthal component of the orientation of orbital in space. Electrons in a particular subshell are defined by values of . The value of ml can range from - to +, including zero. Thus the s, p, d, and f subshells contain 1, 3, 5, and 7 orbitals each, with values of m within the ranges 0, ±1, ±2, ±3 respectively. Each of these orbitals can accommodate up to two electrons, forming the basis of the periodic table.

Valence electron

In chemistry and physics, a valence electron is an outer shell electron that is associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair.

In quantum mechanics, the term symbol is an abbreviated description of the (total) angular momentum quantum numbers in a multi-electron atom. Each energy level of an atom with a given electron configuration is described by not only the electron configuration but also its own term symbol, as the energy level also depends on the total angular momentum including spin. The usual atomic term symbols assume LS coupling. The ground state term symbol is predicted by Hund's rules.

Aufbau principle

The aufbau principle, from the German Aufbauprinzip, also called the aufbau rule, states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s subshell is filled before the 2s subshell is occupied. In this way, the electrons of an atom or ion form the most stable electron configuration possible. An example is the configuration 1s2 2s2 2p6 3s2 3p3 for the phosphorus atom, meaning that the 1s subshell has 2 electrons, and so on.

In atomic physics, Hund's rules refers to a set of rules that German physicist Friedrich Hund formulated around 1927, which are used to determine the term symbol that corresponds to the ground state of a multi-electron atom. The first rule is especially important in chemistry, where it is often referred to simply as Hund's Rule.

Core electrons are the electrons in an atom that are not valence electrons and do not participate in chemical bonding. The nucleus and the core electrons of an atom form the atomic core. Core electrons are tightly bound to the nucleus. Therefore, unlike valence electrons, core electrons play a secondary role in chemical bonding and reactions by screening the positive charge of the atomic nucleus from the valence electrons.

A block of the periodic table is a set of elements unified by the orbitals their valence electrons or vacancies lie in. The term appears to have been first used by Charles Janet. Each block is named after its characteristic orbital: s-block, p-block, d-block, and f-block.

References

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  2. 1 2 Electron Subshells. Corrosion Source.
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