In p-block elements the last electron enters the outermost p orbital. As we know that the number of p orbitals is three and, therefore, the maximum number of electrons that can be accommodated in a set of p orbitals is six.
Consequently there are six groups of p–block elements in the periodic table numbering from 13 to 18. Boron, carbon, nitrogen, oxygen, fluorine and helium head the groups.
Their valence shell electronic configuration is ns2np1-6(except for He). The inner core of the electronic configuration may, however, differ.
The difference in inner core of elements greatly influences their physical properties (such as atomic and ionic radii, ionisation enthalpy, etc.) as well as chemical properties. Consequently, a lot of variation in properties of elements in a group of p-block is observed.
The maximum oxidation state shown by a p-block element is equal to the total number of valence electrons (i.e., the sum of the sand p-electrons).
Clearly, the number of possible oxidation states increases towards the right of the periodic table. In addition to this so called group oxidation state, p-block elements may show other oxidation states which normally, but not necessarily, differ from the total number of valence electrons by unit of two.
The important oxidation states exhibited by p-block elements are shown in Table. In boron, carbon and nitrogen families the group oxidation state is the most stable state for the lighter elements in the group.
However, the oxidation state two unit less than the group oxidation state becomes progressively more stable for the heavier elements in each group.
The occurrence of oxidation states two unit less than the group oxidation states are sometime attributed to the ‘inert pair effect’.
The relative stabilities of these two oxidation states – group oxidation state and two unit less than the group oxidation state – may vary from group to group and will be discussed at appropriate places.
It is interesting to note that the non-metals and metalloids exist only in the p-block of the periodic table. The non-metallic character of elements decreases down the group.
In fact the heaviest element in each p-block group is the most metallic in nature. This change from nonmetallic to metallic character brings diversity in the chemistry of these elements depending on the group to which they belong. In general, non-metals have higher ionisation enthalpies and higher electronegativities than the metals.
Hence, in contrast to metals which readily form cations, non-metals readily form anions. The compounds formed by highly reactive non-metals with highly reactive metals are generally ionic because of large differences in their electronegativities.
On the other hand, compounds formed between non-metals themselves are largely covalent in character because of small differences in their electronegativities.
The change of non-metallic to metallic character can be best illustrated by the nature of oxides they form. The non-metal oxides are acidic or neutral whereas metal oxides are basic in nature.
The first member of p-block differs from the remaining members of their corresponding group in two major respects. First is the size and all other properties which depend on size.
Thus, the lightest p-block elements show the same kind of differences as the lightest s-block elements, lithium and beryllium.
The second important difference, which applies only to the p-block elements, arises from the effect of dorbitals in the valence shell of heavier elements (starting from the third period onwards) and their lack in second period elements.
The second period elements of p-groups starting from boron are restricted to a maximum covalence of four (using 2s and three 2p orbitals).
In contrast, the third period elements of p-groups with the electronic configuration 3s23pn have the vacant 3d orbitals lying between the 3p and the 4s levels of energy. Using these d-orbitals the third period elements can expand their covalence above four.
For example, while boron forms only [BF4]–, aluminium gives [AlF6]3– ion. The presence of these d-orbitals influences the chemistry of the heavier elements in a number of other ways. The combined effect of size and availability of d orbitals considerably influences the ability of these elements to form bonds.
The first member of a group differs from the heavier members in its ability to form p - p multiple bonds to itself and to other second row elements (e.g., C=O, C=N, CN, N=O). This type of ‑ - bonding is not particularly strong for the heavier p-block elements.
The heavier elements do form ‑ bonds but this involves d orbitals (d‑ – p‑ or d‑ –d‑ ). As the d orbitals are of higher energy than the p orbitals, they contribute less to the overall stability of molecules than does p‑ - p‑ bonding of the second row elements.
However, the coordination number in species of heavier elements may be higher than for the first element in the same oxidation state.
For example, in +5 oxidation state both N and P form oxoanions : NO3– (three-coordination with ‑ – bond involving one nitrogen p-orbital) and PO4 3− (four-coordination involving s, p and d orbitals contributing to the ‑ – bond). In this unit we will study the chemistry of group 13 and 14 elements of the periodic table.