Question

How to study the Electronic Configuration of all the atoms ?

Answer 1

The first 20 I can explain as they are easy - Firstly you just look on the periodic table in which period (horizontal row) and group (vertical column) the element is in. The period represents how may energy levels is in its atoms, and the group represents the number of valence electrons (outermost electrons) present. Now from the theory it is true that in the first energy level there is only 1 s orbital, in the 2nd and 3rd energy levels there is 1 s and 3 p orbitals. Now you must also remember the rules for ho electrons fill up - from lowest energy levels first closest to the nucleus, and a maximum of 2 electrons can be in a single orbital provided they spin in opposite directions. Vacant orbitals are occupies singly first. So for example if you take a random element like say Sulphur (S). It is situated in period 3, group 6 on the periodic table. Hence it has 3 main energy levels, and 6 valence electrons in the outermost (3rd) energy level. It has atomic number 16 and hence a neutral S atom will also have 16 electrons. So putting all of this together we get that its electron configuration is `1s^(2)2s^(2) 2p^(6) 3s^(2) 3p^(4)` and hence of the 6 valence electrons in the 3rd energy level, 2 are paired and 2 unpaired, hence it has a valency of 2 and can form 2 bonds.
For the elements above Z=20, we get d and f orbitals and in addition, some elements like carbon can occur in a hybridized state and electrons may rise into higher energy levels to form additional bonds.