The electron configuration of an element is a list of the atomic orbitals which are occupied by electrons, and how many electrons are in each of those orbitals. The rules for writing electron configurations are known as the aufbau principle (German for "building up"):
- Each electron that is added to an atom is placed in the lowest-energy orbital that is available. The orbitals are filled in the order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
Each orbital can hold no more than two electrons. Two electrons in the same orbital must have opposite spins (the Pauli exclusion principle).
If two or more orbitals are available at the same energy level (degenerate orbitals), one electron is placed in each orbital until the available orbitals are occupied by one electron; any additional electrons are placed in the half-filled orbitals.
Electron configurations are written as a list of orbitals which are occupied, followed by a superscript to indicate how many electrons are in those orbitals.
H 1s1 He 1s2 Li 1s2 2s1 Be 1s2 2s2 B 1s2 2s2 2p1 C 1s2 2s2 2p2 N 1s2 2s2 2p3 O 1s2 2s2 2p3 F 1s2 2s2 2p4 Ne 1s2 2s2 2p5 Na 1s2 2s2 2p6 3s1
Electron configurations in which all of the electrons are in their lowest-energy configurations are known as ground state configurations. If an electron absorbs energy, it can move into a higher-energy orbital, producing an excited state configuration.
For atoms with a large number of electrons, the complete electron can be very cumbersome, and not very informative. For instance, the complete configuration of the element radium is
Ra: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 6s2 6p6 7s2
(With a description like that, you'd be radioactive too!) Since everything up to the 6p6 is the same electron configuration as the noble gas radon, the configuration can be abbreviated as
Ra: [Rn] 7s2
Abbreviated electron configurations are always based on the nearest preceding noble gas.
Electron configurations can be written directly from the periodic table, without having to memorize the aufbau scheme, using the following patterns:
Half-filled and filled subshells are especially stable, leading to some anomalous electron configurations:
Predicted configuration Actual configuration Cr [Ar] 3d4 4s2 [Ar] 3d5 4s1 Cu [Ar] 3d9 4s2 [Ar] 3d10 4s1 Ag [Kr] 4d9 5s2 [Kr] 4d10 5s1 Au [Xe] 4f14 5d9 6s2 [Xe] 4f14 5d10 6s1
In the case of chromium, an electron from the 4s orbital moves into a 3d orbital, allowing each of the five 3d orbitals to have one electron, making a half-filled set of orbitals. In the case of copper, silver and gold, an electron from the highest-occupied s orbital moves into the d orbitals, thus filling the d subshell. Many anomalous electron configurations occur in the heavier transition metals and inner transition metals, where the energy differences between the s, d, and f subshells is very small.