(i) Electronic configuration: The valence shell electronic configuration of these elements is ns² np³. The s orbital in these elements is completely filled and p orbitals are half filled, making their electronic configuration extra stable.

(ii) Atomic Size: Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and f or f orbitals in heavier members. 

(iii)Oxidation State: The common oxidation states of these elements are –3, + 3 and + 5. The tendency to exhibit –3 oxidation state decreases down the group due to increase in size and metallic group. In the last member of the group, bismuth hardly forms any compound in –3 oxidation state. The stability of + 5 oxidation state decreases down the group. The stability of + 5 oxidation state decreases and that of + 3 state increases (due to invert pair effect) down the group. Nitrogen exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. Phosphorus also shows +1 and + 4 oxidation states in some oxo acids.

(iv) Ionization enthalpy: Ionization enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half filled p orbitals electronic configuration and smaller size, the ionization enthalpy of the group 15 elements is much greater than that of group 14 elements in the corresponding periods. The order of successive ionization emthalpies are expected as ΔH₁, < ΔH₂ < ΔH₃.

(v) Electronegativity: The electronegativity value, in general, decreases down the group with increasing atomic size. However, amongst the heavier elements, the different is not that much pronounced.