Periodic table/Electron configurations

This version of the periodic table shows for each element its electron configuration, in other words how electrons are distributed over various shells.

For each row of elements only new shells are displayed, all shells shown on previous rows are fully occupied (to obtain the full set add the configurations of the most rightmost element, the noble gases, in all rows above).

Example: the full electron configuration of Bromine (symbol Br) is 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁵. Often a shorthand notation is used: all shells that already occurred on previous rows in the system are left out and replaced by the symbol of the noble gas in the row just above the element being described. The electron configuration of bromine then is coded as follows: [Ar] 4s²3d¹⁰4p⁵.

With increasing atomic number shells and sub shells are filled in a fairly consistent way: first 1s, then 2s, 2p, 3s and 3p, but 4s is filled before 3d. Even within this exception there is another irregularity: for instance, Copper has configuration 3d¹⁰4s¹ in stead of 3d⁹4s². The reason for these irregularities is that besides the attraction between the positive atomic nucleus and the electrons the latter repel each other as well. As a result of this many-body problem energy levels of the various orbits interfere somewhat like a rope ladder: put your foot on it and all distances between steps change.

The same problem occurs during ion formation. For instance, when Manganese (symbol Mn) (3d⁵4s²) dissolves in water it forms Mn²⁺ ions which have a 3d⁵ configuration. Although 4s² shells are filled before 3d still the former give up electrons during oxidation into Mn²⁺. Yet another factor comes into play: at 3d⁵ the d-subshell is exactly half full, which provides extra stability, provided that all spins are set paralel. The spin pair energy is a rather small term but it can be decisive in rendering a certain configuration stable.

(..) = this element does not occur on earth