The rules in filling orbitals:

There are two rules in filling orbitals, the first rule is for hydrogen atom, the second rule is for the other atoms. For hydrogen atom, the energy of the electron is determined solely by its principal quantum number. Thus, the energies of hydrogen orbitals increase as follows:

1s< 2s= 2p < 3s = 3p = 3d< 4s = 4p = 4d = 4 f <…………

The electron in “1s” is in the ground state and is in the most stable condition. For 2s and 2p orbitals, the orbitals have the same energies, and the electron is in its excited state.

For other atoms, the electron distribution depends on the angular momentum quantum number and the principal quantum number. For many atoms, the 3d energy level is close to the 4s energy level. Yet, the total energy of an atom is lower when 4s subshell is filled first before 3d subshell. The reason is that the total energy of an atom depends on the sum of orbital energies and the repulsion between the electrons (each orbital can take up to 2 electrons).

For 3p and 4s, electrons would prefer to fill 3p orbital before 4s. If the 3p orbital is filled first, the repulsion between the electrons will be minimum comparing to 4s because the two electrons will fill 3p_x and 3p_y and will be unpaired. While filling 4s would mean that electrons have to be paired. For that, the repulsion between electrons in 3p orbital is less than 4s orbital.