The word "Aufbau" comes from the German meaning "construction". It is how the electrons fill up the available atomic orbitals in order of increasing energy, according to certain rules.
Syllabus reference S1.3.5Structure 1.3.5 - Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.
- Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.
- Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.
Guidance
- Full electron configurations and condensed electron configurations using the noble gas core should be covered
- Orbital diagrams, i.e. arrow-in-box diagrams, should be used to represent the filling and relative energy of orbitals.
- The electron configurations of Cr and Cu as exceptions should be covered.
Tools and links
Atomic orbitals
Each energy level is split into sub-levels (except energy level 1). The sub-levels in turn, contain orbitals that can hold a maximum of two electrons per orbital.
- Level 1 contains only 1 's' orbital
- Level 2 contains 1 's' and three 'p' orbitals
- Level 3 contains 1 s, three p and five 'd' orbitals
- Level 4 contains one 's', three 'p', five 'd' and seven 'f' orbitals
As stated above the electrons fill up the orbitals in order of increasing energy from the lowest energy orbitals upwards. This process is subject to certain 'rules
Pauli's exclusion principle
This states that no two electrons can be identical within an atom. Simply
stated, it means that only two electrons can fit into each atomic orbital
and they must have opposite spins. By convention, we say that one electron
spins up, 
and the other down 
represented by up and down arrows.
Hund's rule
This states that electrons entering orbitals that have the same energy (degenerate orbitals, represented by boxes on the same level) must be filled by parallel electrons (unpaired electrons), before the electrons become paired up.
Example: In the electronic configuration of carbon 1s2 2s2 2p2, the electrons fill up in the following way:
Anomalous configurations
The Aufbau principle works fairly well for the first 38 elements, but after that it starts to break down. Even so, there are two configurations that do not seem to fit into the pattern. These are chromium and copper.
Chromium (24 electrons) has an expected configuration of [Ar] 4s2 3d4. However, the actual configuration is [Ar] 4s1 3d5.
This is explained by suggesting that there is some energetic advantage to the atom to have a half-full set of 'd' orbitals, and that this is enough to cause one of the 4s orbital electrons to occupy the last orbital in the 3d series.
Copper (29 electrons) also has an anomalous configuration with the expected [Ar] 4s2 3d9 giving way to [Ar] 4s1 3d10.
Once again, this is explained by the extra stability due to a full set of 3d orbitals, providing the incentive for the 4s electron to be housed in the last 3d orbital.
Example: Which of the following atoms has/have one or more unpaired electrons?
- I. Iron
- II. Copper
- III. Zinc
To answer the question it is necessary to look at the electronic configuration of each atom.
- Iron [Ar] 4s2 3d6
- Copper [Ar] 4s1 3d10
- Zinc [Ar] 4s2 3d10
Iron has 6 'd' electrons to fit into 5 'd' orbitals. As the first five must enter with parallel spin it has 4 unpaired electrons
Copper has a single unpaired electron in the 4s orbital and all the '3d' orbitals are full
Zinc has a full set of 4s and 3d orbitals
Therefore, of the three elements, only iron and copper have unpaired electrons.
Electronic configuration of ions
Ions are formed from atoms by the addition or removal of electrons depending on whether the atom is a metal or a non-metal.
Metals lose electrons forming positive ions. The number of electrons lost depends on the metal atoms. Group 1, 2 and 3 elements lose 1, 2 and 3 electrons respectively to give a noble gas configuration.
The 'd' block (transition) metals have variable oxidation states and may lose a variable number of electrons. The first electrons lost by the first row 'd' block metals are the 4s electrons. After these have been removed the '3d' electrons are successively removed until the required ion is obtained.
Example: Show the electronic configuration of iron(II) and iron(III) ions.
iron(II) has a charge of 2+ and has consequently lost 2 electrons. Iron (III) has a charge of 3+ and has lost 3 electrons.
Iron (at no. 26) has an electronic configuration = 1s2 2s2 2p6 3s2 3p6 4s2 3d6
iron(II) has a configuration of 1s2 2s2 2p6 3s2 3p6 4s0 3d6
iron(III) has a configuration of 1s2 2s2 2p6 3s2 3p6 4s0 3d5