Electron configuration
In atomic physics and quantum chemistry, the electron configuration is the arrangement of
electrons in an
atom,
molecule or other body. The electrons occupy specific probability regions , whose shapes and electron capacity vary. There can be
atomic orbitals,
molecular orbitals, or other types of electron orbitals. Electrons are able to jump from one energy level to another by emission of a quantum of energy, in the form of a
photon. The properties of electron configuration in atoms form the basis for
covalent bonding.
Orbital shapes and electron capacity are denoted by the letters s, p, d, f and the as-of-yet unseen g, h and i.
Encyclopedia
In atomic physics and quantum chemistry, the
electron configuration is the arrangement of
electrons in an
atom,
molecule or other body. The electrons occupy specific probability regions , whose shapes and electron capacity vary. There can be
atomic orbitals,
molecular orbitals, or other types of electron orbitals. Electrons are able to jump from one energy level to another by emission of a quantum of energy, in the form of a
photon. The properties of electron configuration in atoms form the basis for
covalent bonding.
Orbital shapes and electron capacity are denoted by the letters s, p, d, f and the as-of-yet unseen g, h and i. The energy of an orbital is shown by a whole number next to the letter.
The subshell labels s, p, d, and f originate from a now-discredited system of categorizing
spectral lines as
sharp,
principal,
diffuse, or
fundamental, based on their observed fine structure. When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designation
g was derived by following alphabetical order. Shells with more than five subshells are theoretically permissible, but this covers all discovered elements.
For
mnemonic reasons, some call the s and p orbitals
spherical and
peripheral.
Electron configuration in atoms
The discussion below presumes knowledge of material contained at Atomic orbital.Summary of the quantum numbers
The state of an electron in an atom is given by four quantum numbers. Three of these are integers and are properties of the
atomic orbital in which it sits .
| number | denoted | allowed range | represents |
|---|
| principal quantum number | n | integer, 1 or more | partly the overall energy of the orbital, and by extension its general distance from the nucleus |
| azimuthal quantum number | l | integer, 0 to n-1 | the orbital's angular momentum, also seen as the number of nodes in the density plot |
| magnetic quantum number | m | integer, -l to +l | determines energy shift of an atomic orbital due to external magnetic field . |
| spin quantum number | s | ½ or -½ Spin is an intrinsic property of the electron and independent of the other numbers. s and l in part determine the electron's magnetic dipole moment. |
No two electrons in one atom can have the same set of these four quantum numbers .
Shells and subshells
Shells and subshells are defined by the quantum numbers, NOT by the distance of its electrons from the nucleus. In large atoms, shells above the second shell overlap .
States with the same value of
n are related, and said to lie within the same
electron shell.
States with the same value of
n and also
l are said to lie within the same
electron subshell.
If the states also share the same value of
m, they are said to lie in the same
atomic orbital.
Because electrons have only two possible spin states, an atomic orbital cannot contain more than two electrons .
A subshell can contain up to 4
l+2 electrons; a shell can contain up to 2
n² electrons.
Worked example
Here is the electron configuration for a filled fifth shell:
| Shell | Subshell | Orbitals | | Electrons |
| n = 5 | l = 0 | m = 0 | ? 1 type s orbital | ? max 2 electrons |
| | l = 1 | m = -1, 0, +1 | ? 3 type p orbitals | ? max 6 electrons |
| | l = 2 | m = -2, -1, 0, +1, +2 | ? 5 type d orbitals | ? max 10 electrons |
| | l = 3 | m = -3, -2, -1, 0, +1, +2, +3 | ? 7 type f orbitals | ? max 14 electrons |
| | l = 4 | m = -4, -3 -2, -1, 0, +1, +2, +3, +4 | ? 9 type g orbitals | ? max 18 electrons |
| | | | | Total: max 50 electrons |
|
This information can be written as 5s
2 5p
6 5d
10 5f
14 5g
18 .
Notation
Physicists and chemists use a standard notation to describe atomic electron configurations. In this notation, a subshell is written in the form n
xe, where n is the shell number,
x is the subshell label and e is the number of electrons in the subshell. An atom's subshells are written in order of increasing energy - in other words, the sequence in which they are filled .
For instance, ground-state
hydrogen has one electron in the s subshell of the first shell, so its configuration is written 1s
1.
Lithium has two electrons in the 1s subshell and one in the 2
s subshell, so its ground-state configuration is written 1
s2 2
s1.
Phosphorus , is as follows: 1
s2 2
s2 2
p6 3
s2 3
p3.
For atoms with many electrons, this notation can become lengthy. It is often abbreviated by noting that the first few subshells are identical to those of one or another noble gas. Phosphorus, for instance, differs from
neon only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne]3
s2 3
p3.
An even simpler version is simply to quote the number of electrons in each shell, e.g. : 2-8-5.
Aufbau principle
In the ground state of an atom , the electron configuration generally follows the
Aufbau principle. According to this principle, electrons enter into states in order of the states' increasing energy; i.e., the first electron goes into the lowest-energy state, the second into the next lowest, and so on. The order in which the states are filled is as follows:
| | | | | |
|---|
| 1 | 1 |
|---|
| 2 | 2 | 3 |
|---|
| 3 | 4 | 5 | 7 |
|---|
| 4 | 6 | 8 | 10 | 13 |
|---|
| 5 | 9 | 11 | 14 | 17 | 21 |
|---|
| 6 | 12 | 15 | 18 | 22 |
|---|
| 7 | 16 | 19 | 23 |
|---|
| 8 | 20 | 24 |
|---|
The order of increasing energy of the subshells can be constructed by going through downward-leftward diagonals of the table above , going from the topmost diagonals to the bottom. The first diagonal goes through 1s; the second diagonal goes through 2s; the third goes through 2p and 3s; the fourth goes through 3p and 4s; the fifth goes through 3d, 4p, and 5s; and so on. In general, a subshell that is not "s" is always followed by a "lower" subshell of the next shell; e.g. 2p is followed by 3s; 3d is followed by 4p, which is followed by 5s, 4f is followed by 5d, which is followed by 6p, and then 7s. This explains the ordering of the blocks in the periodic table.
A pair of electrons with identical spins has slightly less energy than a pair of electrons with opposite spins. Since two electrons in the same orbital must have opposite spins, this causes electrons to prefer to occupy different orbitals. This preference manifests itself if a subshell with is less than full. For instance, if a
p subshell contains four electrons, two electrons will be forced to occupy one orbital, but the other two electrons will occupy
both of the other orbitals, and their spins will be equal. This phenomenon is called
Hund's rule.
The Aufbau principle can be applied, in a modified form, to the
protons and
neutrons in the
atomic nucleus .
Exceptions in 3d, 4d, 5d
A
d subshell that is half-filled or full is more stable than the
s subshell of the next shell. This is the case because it takes less energy to maintain an electron in a half-filled
d subshell than a filled
s subshell. For instance,
copper has a configuration of [Ar]4s
1 3d
10, not [Ar]4s
2 3d
9 as one would expect by the Aufbau principle. Likewise,
chromium has a configuration of [Ar]4s
1 3d
5, not [Ar]4s
2 3d
4.
| Element | Z | Electron configuration | Short electron conf. |
| Scandium | 21 | 1s2 2s2 2p6 3s2 3p6 4s2 3d1
| [Ar] 4s2 3d1 |
| Titanium | 22 | 1s2 2s2 2p6 3s2 3p6 4s2 3d2
| [Ar] 4s2 3d2 |
| Vanadium | 23 | 1s2 2s2 2p6 3s2 3p6 4s2 3d3
| [Ar] 4s2 3d3 |
| Chromium | 24 | 1s2 2s2 2p6 3s2 3p6 4s1 3d5
| [Ar] 4s1 3d5 |
| Manganese | 25 | 1s2 2s2 2p6 3s2 3p6 4s2 3d5
| [Ar] 4s2 3d5 |
| Iron | 26 | 1s2 2s2 2p6 3s2 3p6 4s2 3d6
| [Ar] 4s2 3d6 |
| Cobalt | 27 | 1s2 2s2 2p6 3s2 3p6 4s2 3d7
| [Ar] 4s2 3d7 |
| Nickel | 28 | 1s2 2s2 2p6 3s2 3p6 4s2 3d8
| [Ar] 4s2 3d8 |
| Copper | 29 | 1s2 2s2 2p6 3s2 3p6 4s1 3d10
| [Ar] 4s1 3d10 |
| Zinc | 30 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10
| [Ar] 4s2 3d10 |
| Gallium | 31 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
| [Ar] 4s2 3d10 4p1 |
|
This can be most easily understood by stepping through the electron configuration shown at
Period 5
th has more exceptions:
| Element | Z | Electron configuration | Short electron conf. |
| Yttrium | 39 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d1
| [Kr] 5s2 4d1 |
| Zirconium | 40 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2
| [Kr] 5s2 4d2 |
| Niobium | 41 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d4
| [Kr] 5s1 4d4 |
| Molybdenum | 42 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d5
| [Kr] 5s1 4d5 |
| Technetium | 43 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5
| [Kr] 5s2 4d5 |
| Ruthenium | 44 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d7
| [Kr] 5s1 4d7 |
| Rhodium | 45 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d8
| [Kr] 5s1 4d8 |
| Palladium | 46 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 4d10
| [Kr] 4d10 |
| Silver | 47 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10
| [Kr] 5s1 4d10 |
| Cadmium | 48 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
| [Kr] 5s2 4d10 |
| Indium | 49 | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1
| [Kr] 5s2 4d10 5p1 |
|
This can be seen by stepping through the electron configuration shown at
| Element | Z | Short electron conf. |
| Iridium | 77
| [Xe] 6s2 4f14 5d7 |
| Platinum | 78
| [Xe] 6s1 4f14 5d9 |
| Gold | 79
| [Xe] 6s1 4f14 5d10 |
| Mercury | 80
| [Xe] 6s2 4f14 5d10 |
| Thallium | 81
| [Xe] 6s2 4f14 5d10 6p1 |
|
This can be seen by stepping through the electron configuration shown at
Relation to the structure of the periodic table
Electron configuration is intimately related to the structure of the periodic table. The chemical properties of an atom are largely determined by the arrangement of the electrons in its outermost
shell .
Electron configuration in molecules
In molecules, the situation becomes more complex, as each molecule has a different orbital structure. See the
molecular orbital article and the
linear combination of atomic orbitals method for an introduction and the
computational chemistry article for more advanced discussions.
Electron configuration in solids
In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend together into continuous ranges of possible states . The notion of electron configuration ceases to be relevant, and yields to
band theory.
See also
- Atomic electron configuration table
- Periodic table
- Atomic orbital
- Energy level
- Molecular term symbol
- HOMO/LUMO
- Atompaw Software package for electron configuration calculations
- Pwpaw Software package for electron configuration calculations
