Electron shell

From Wikipedia, the free encyclopedia - View original article

 
Jump to: navigation, search
Periodic table with electron shells

In chemistry and atomic physics, an electron shell, also called a principal energy level 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 1st shell can hold up to two electrons, the 2nd shell can hold up to eight (2 + 6) electrons, the 3rd 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 2n2 electrons.[1] 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.[2]

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.

History[edit]

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. 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[edit]

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[edit]

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, although occupied in actinides, is not filled in any element occurring naturally.[2] The various possible subshells are shown in the following table:

Subshell labelMax 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)[4]

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,[5] 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[edit]

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

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; that's why nth shell can hold up to 2n2 electrons.[1]

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).

Valence shells[edit]

Main article: Valence electron

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 (alkali metals) or just missing one electron from having a complete shell (halogens) are the most reactive.[8]

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 highest energy. For the transition elements, the partially filled (n − 1)d energy level is very close in energy to the ns level[9] and hence the d electrons in transition metals behave as valence electrons although they are not in the so-called valence shell.

List of elements with electrons per shell[edit]

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   electric 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).

ZElementNo. of electrons/shellGroup
1Hydrogen11
2Helium218
3Lithium2, 11
4Beryllium2, 22
5Boron2, 313
6Carbon2, 414
7Nitrogen2, 515
8Oxygen2, 616
9Fluorine2, 717
10Neon2, 818
11Sodium2, 8, 11
12Magnesium2, 8, 22
13Aluminium2, 8, 313
14Silicon2, 8, 414
15Phosphorus2, 8, 515
16Sulfur2, 8, 616
17Chlorine2, 8, 717
18Argon2, 8, 818
19Potassium2, 8, 8, 11
20Calcium2, 8, 8, 22
21Scandium2, 8, 9, 23
22Titanium2, 8, 10, 24
23Vanadium2, 8, 11, 25
24Chromium2, 8, 13, 16
25Manganese2, 8, 13, 27
26Iron2, 8, 14, 28
27Cobalt2, 8, 15, 29
28Nickel2, 8, 16, 210
29Copper2, 8, 18, 111
30Zinc2, 8, 18, 212
31Gallium2, 8, 18, 313
32Germanium2, 8, 18, 414
33Arsenic2, 8, 18, 515
34Selenium2, 8, 18, 616
35Bromine2, 8, 18, 717
36Krypton2, 8, 18, 818
37Rubidium2, 8, 18, 8, 11
38Strontium2, 8, 18, 8, 22
39Yttrium2, 8, 18, 9, 23
40Zirconium2, 8, 18, 10, 24
41Niobium2, 8, 18, 12, 15
42Molybdenum2, 8, 18, 13, 16
43Technetium2, 8, 18, 13, 27
44Ruthenium2, 8, 18, 15, 18
45Rhodium2, 8, 18, 16, 19
46Palladium2, 8, 18, 1810
47Silver2, 8, 18, 18, 111
48Cadmium2, 8, 18, 18, 212
49Indium2, 8, 18, 18, 313
50Tin2, 8, 18, 18, 414
51Antimony2, 8, 18, 18, 515
52Tellurium2, 8, 18, 18, 616
53Iodine2, 8, 18, 18, 717
54Xenon2, 8, 18, 18, 818
55Caesium2, 8, 18, 18, 8, 11
56Barium2, 8, 18, 18, 8, 22
57Lanthanum2, 8, 18, 18, 9, 2
58Cerium2, 8, 18, 19, 9, 2
59Praseodymium2, 8, 18, 21, 8, 2
60Neodymium2, 8, 18, 22, 8, 2
61Promethium2, 8, 18, 23, 8, 2
62Samarium2, 8, 18, 24, 8, 2
63Europium2, 8, 18, 25, 8, 2
64Gadolinium2, 8, 18, 25, 9, 2
65Terbium2, 8, 18, 27, 8, 2
66Dysprosium2, 8, 18, 28, 8, 2
67Holmium2, 8, 18, 29, 8, 2
68Erbium2, 8, 18, 30, 8, 2
69Thulium2, 8, 18, 31, 8, 2
70Ytterbium2, 8, 18, 32, 8, 2
71Lutetium2, 8, 18, 32, 9, 23
72Hafnium2, 8, 18, 32, 10, 24
73Tantalum2, 8, 18, 32, 11, 25
74Tungsten2, 8, 18, 32, 12, 26
75Rhenium2, 8, 18, 32, 13, 27
76Osmium2, 8, 18, 32, 14, 28
77Iridium2, 8, 18, 32, 15, 29
78Platinum2, 8, 18, 32, 17, 110
79Gold2, 8, 18, 32, 18, 111
80Mercury2, 8, 18, 32, 18, 212
81Thallium2, 8, 18, 32, 18, 313
82Lead2, 8, 18, 32, 18, 414
83Bismuth2, 8, 18, 32, 18, 515
84Polonium2, 8, 18, 32, 18, 616
85Astatine2, 8, 18, 32, 18, 717
86Radon2, 8, 18, 32, 18, 818
87Francium2, 8, 18, 32, 18, 8, 11
88Radium2, 8, 18, 32, 18, 8, 22
89Actinium2, 8, 18, 32, 18, 9, 2
90Thorium2, 8, 18, 32, 18, 10, 2
91Protactinium2, 8, 18, 32, 20, 9, 2
92Uranium2, 8, 18, 32, 21, 9, 2
93Neptunium2, 8, 18, 32, 22, 9, 2
94Plutonium2, 8, 18, 32, 24, 8, 2
95Americium2, 8, 18, 32, 25, 8, 2
96Curium2, 8, 18, 32, 25, 9, 2
97Berkelium2, 8, 18, 32, 27, 8, 2
98Californium2, 8, 18, 32, 28, 8, 2
99Einsteinium2, 8, 18, 32, 29, 8, 2
100Fermium2, 8, 18, 32, 30, 8, 2
101Mendelevium2, 8, 18, 32, 31, 8, 2
102Nobelium2, 8, 18, 32, 32, 8, 2
103Lawrencium2, 8, 18, 32, 32, 10, 1 (?)3
104Rutherfordium2, 8, 18, 32, 32, 10, 2 (?)4
105Dubnium2, 8, 18, 32, 32, 11, 2 (?)5
106Seaborgium2, 8, 18, 32, 32, 12, 2 (?)6
107Bohrium2, 8, 18, 32, 32, 13, 2 (?)7
108Hassium2, 8, 18, 32, 32, 14, 2 (?)8
109Meitnerium2, 8, 18, 32, 32, 15, 2 (?)9
110Darmstadtium2, 8, 18, 32, 32, 16, 2 (?)10
111Roentgenium2, 8, 18, 32, 32, 18, 1 (?)11
112Copernicium2, 8, 18, 32, 32, 18, 2 (?)12
113Ununtrium2, 8, 18, 32, 32, 18, 3 (?)13
114Flerovium2, 8, 18, 32, 32, 18, 4 (?)14
115Ununpentium2, 8, 18, 32, 32, 18, 5 (?)15
116Livermorium2, 8, 18, 32, 32, 18, 6 (?)16
117Ununseptium2, 8, 18, 32, 32, 18, 7 (?)17
118Ununoctium2, 8, 18, 32, 32, 18, 8 (?)18

See also[edit]

References[edit]

  1. ^ a b Re: Why do electron shells have set limits ? madsci.org, 17 March 1999, Dan Berger, Faculty Chemistry/Science, Bluffton College
  2. ^ a b Electron Subshells. Corrosion Source. Retrieved on 1 December 2011.
  3. ^ Donald Sadoway, Introduction to Solid State Chemistry, Lecture 5
  4. ^ Jue, T. (2009). "Quantum Mechanic Basic to Biophysical Methods". Fundamental Concepts in Biophysics. Berlin: Springer. p. 33. ISBN 1-58829-973-2. 
  5. ^ 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.
  6. ^ Orbitals. Chem4Kids. Retrieved on 1 December 2011.
  7. ^ Electron & Shell Configuration. Chemistry.patent-invent.com. Retrieved on 1 December 2011.
  8. ^ Chemical Reactions. Vision Learning (26 July 2011). Retrieved on 1 December 2011.
  9. ^ THE ORDER OF FILLING 3d AND 4s ORBITALS. chemguide.co.uk