(Kr). In the next eighteen elements from rubidium (Rb) to xenon (Xe), the pattern of filling the s , d and p orbitals are similar to that of s , d and p orbitals as discussed above. Then comes the turn of the s orbital. In caesium (Cs) and the barium (Ba), this orbital contains one and two electrons, respectively.
Then from lanthanum (La) to mercury (Hg), the filling up of electrons takes place in f and d orbitals. After this, filling of p , then s and finally f and d orbitals takes place. The elements after uranium (U) are all short-lived and all of them are produced artificially. The electronic configurations of the known elements (as determined by spectroscopic methods) are tabulated in Table .
(page ). One may ask what is the utility of knowing the electron configuration? The modern approach to the chemistry, infact, depends almost entirely on electronic distribution to understand and explain chemical behaviour. For example, questions like why two or more atoms combine to form molecules, why some elements are metals while others are non- metals, why elements like helium and argon are not reactive but elements like the halogens are reactive, find simple explanation from the electronic configuration.
These questions have no answer in the Daltonian model of atom. A detailed understanding of the electronic structure of atom is, therefore, very essential for getting an insight into the various aspects of modern chemical knowledge. . .
Stability of Completely Filled and Half Filled Subshells The ground state electronic configuration of the atom of an element always corresponds to the state of the lowest total electronic energy. The electronic configurations of most of the atoms follow the basic rules given in Section . . .
However, in certain elements such as Cu, or Cr, where the two subshells ( s and d ) differ slightly in their energies, an electron shifts from a subshell of lower energy ( s ) to a subshell of higher energy ( d ), provided such a