It has been observed that Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, and Fluorine have slightly different periodic properties than the rest of the elements belonging to Groups 1, 2, 13-17.
Work Step by Step
For example, Lithium and Beryllium form covalent compounds, whereas the rest of the members of Groups 1 and 2 form ionic compounds. Also, the oxide that is formed by Beryllium when it reacts with Oxygen is amphoteric in nature, unlike other Group 2 elements that form basic oxides. Yet another example is that of Carbon which can form stable multiple bonds, whereas Si=Si double bonds are not very common. So, it has clearly been established that the second-period elements are different. In fact, they display periodic properties that are similar to the second element of the next group (i.e. Lithium is similar to Magnesium and Beryllium to Aluminium) or in other words, they have a diagonal relationship. The reasons for differences in periodic properties and hence in chemical behavior are: Small size of these atoms High electronegativity Large charge/radius ratio These elements also have only 4 valence orbitals available (2s and 2p) for bonding as compared to the 9 available (3s, 3p, and 3d) to the other members of the respective groups, so their maximum covalency is 4. (This is why Boron can only form [BF4]– whereas Aluminium can form [AlF6]3-).