Electron shell
Electron shell
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 farther and
farther from the nucleus. The shell
letters K, L, M, ... used in the X-ray
notation 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
Periodic table with electron shells
shell can hold up to 18, and 4th shell
can hold up to 32 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 the valence shell.
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 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.
1
Electron shell
2
Shells
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
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 but the "5f" subshell is not filled in the elements occurring naturally. ;.[1] 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 (theoretically) (next in alphabet after f)[3]
• 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
Electron shell
3
Shell Subshell Subshell
Shell
name name
max
max
electrons electrons
K
1s
2
2
L
2s
2
2+6=8
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
4f
14
M
N
2+6+10
=18
2+6+
+10+14
=32
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
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. For the transition elements, the partially
filled (n-1)d energy level is very close in energy to the ns level[8] and hence the d electrons in transition metals
behave as valence electrons although they are not in the so-called valence shell.
Electron shell
4
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, 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 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, some
atoms have such short half-life that it is impossible to measure their electron configurations).
Z
Element
No. of electrons/shell
Group
1
Hydrogen
1
1
2
Helium
2
18
3
Lithium
2, 1
1
4
Beryllium
2, 2
2
5
Boron
2, 3
13
6
Carbon
2, 4
14
7
Nitrogen
2, 5
15
8
Oxygen
2, 6
16
9
Fluorine
2, 7
17
10
Neon
2, 8
18
11
Sodium
2, 8, 1
1
12
Magnesium
2, 8, 2
2
13
Aluminium
2, 8, 3
13
14
Silicon
2, 8, 4
14
15
Phosphorus
2, 8, 5
15
16
Sulfur
2, 8, 6
16
17
Chlorine
2, 8, 7
17
18
Argon
2, 8, 8
18
19
Potassium
2, 8, 8, 1
1
20
Calcium
2, 8, 8, 2
2
21
Scandium
2, 8, 9, 2
3
22
Titanium
2, 8, 10, 2
4
23
Vanadium
2, 8, 11, 2
5
24
Chromium
2, 8, 13, 1
6
25
Manganese
2, 8, 13, 2
7
26
Iron
2, 8, 14, 2
8
27
Cobalt
2, 8, 15, 2
9
Electron shell
5
28
Nickel
2, 8, 16, 2
10
29
Copper
2, 8, 18, 1
11
30
Zinc
2, 8, 18, 2
12
31
Gallium
2, 8, 18, 3
13
32
Germanium
2, 8, 18, 4
14
33
Arsenic
2, 8, 18, 5
15
34
Selenium
2, 8, 18, 6
16
35
Bromine
2, 8, 18, 7
17
36
Krypton
2, 8, 18, 8
18
37
Rubidium
2, 8, 18, 8, 1
1
38
Strontium
2, 8, 18, 8, 2
2
39
Yttrium
2, 8, 18, 9, 2
3
40
Zirconium
2, 8, 18, 10, 2
4
41
Niobium
2, 8, 18, 12, 1
5
42
Molybdenum
2, 8, 18, 13, 1
6
43
Technetium
2, 8, 18, 13, 2
7
44
Ruthenium
2, 8, 18, 15, 1
8
45
Rhodium
2, 8, 18, 16, 1
9
46
Palladium
2, 8, 18, 18
10
47
Silver
2, 8, 18, 18, 1
11
48
Cadmium
2, 8, 18, 18, 2
12
49
Indium
2, 8, 18, 18, 3
13
50
Tin
2, 8, 18, 18, 4
14
51
Antimony
2, 8, 18, 18, 5
15
52
Tellurium
2, 8, 18, 18, 6
16
53
Iodine
2, 8, 18, 18, 7
17
54
Xenon
2, 8, 18, 18, 8
18
55
Caesium
2, 8, 18, 18, 8, 1
1
56
Barium
2, 8, 18, 18, 8, 2
2
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
Electron shell
6
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, 2
3
72
Hafnium
2, 8, 18, 32, 10, 2
4
73
Tantalum
2, 8, 18, 32, 11, 2
5
74
Tungsten
2, 8, 18, 32, 12, 2
6
75
Rhenium
2, 8, 18, 32, 13, 2
7
76
Osmium
2, 8, 18, 32, 14, 2
8
77
Iridium
2, 8, 18, 32, 15, 2
9
78
Platinum
2, 8, 18, 32, 17, 1
10
79
Gold
2, 8, 18, 32, 18, 1
11
80
Mercury
2, 8, 18, 32, 18, 2
12
81
Thallium
2, 8, 18, 32, 18, 3
13
82
Lead
2, 8, 18, 32, 18, 4
14
83
Bismuth
2, 8, 18, 32, 18, 5
15
84
Polonium
2, 8, 18, 32, 18, 6
16
85
Astatine
2, 8, 18, 32, 18, 7
17
86
Radon
2, 8, 18, 32, 18, 8
18
87
Francium
2, 8, 18, 32, 18, 8, 1
1
88
Radium
2, 8, 18, 32, 18, 8, 2
2
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, 3 (?)
3
104 Rutherfordium 2, 8, 18, 32, 32, 10, 2 (?) 4
105 Dubnium
2, 8, 18, 32, 32, 11, 2 (?) 5
Electron shell
7
106 Seaborgium
2, 8, 18, 32, 32, 12, 2 (?) 6
107 Bohrium
2, 8, 18, 32, 32, 13, 2 (?) 7
108 Hassium
2, 8, 18, 32, 32, 14, 2 (?) 8
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 Ununtrium
2, 8, 18, 32, 32, 18, 3 (?) 13
114 Flerovium
2, 8, 18, 32, 32, 18, 4 (?) 14
115 Ununpentium
2, 8, 18, 32, 32, 18, 5 (?) 15
116 Livermorium
2, 8, 18, 32, 32, 18, 6 (?) 16
117 Ununseptium
2, 8, 18, 32, 32, 18, 7 (?) 17
118 Ununoctium
2, 8, 18, 32, 32, 18, 8 (?) 18
References
[1] Electron Subshells (http:/ / www. corrosionsource. com/ handbook/ periodic/ e_subshells. htm). Corrosion Source. Retrieved on 2011-12-01.
[2] Donald Sadoway, Introduction to Solid State Chemistry, Lecture 5 (http:/ / ocw. mit. edu/ courses/ materials-science-and-engineering/
3-091-introduction-to-solid-state-chemistry-fall-2004/ video-lectures/ embed05/ )
[3] Jue, T. (2009). "Quantum Mechanic Basic to Biophysical Methods" (http:/ / books. google. com/ books?id=-W1bL7s--igC& pg=PA33).
Fundamental Concepts in Biophysics. Berlin: Springer. p. 33. ISBN 1-58829-973-2. .
[4] 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.
[5] Orbitals (http:/ / www. chem4kids. com/ files/ atom_orbital. html). Chem4Kids. Retrieved on 2011-12-01.
[6] Electron & Shell Configuration (http:/ / www. chemistry. patent-invent. com/ chemistry/ electron_configuration. html).
Chemistry.patent-invent.com. Retrieved on 2011-12-01.
[7] Chemical Reactions (http:/ / www. visionlearning. com/ library/ module_viewer. php?mid=54). Vision Learning (2011-07-26). Retrieved on
2011-12-01.
[8] (http:/ / www. chemguide. co. uk/ atoms/ properties/ 3d4sproblem. html)
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