`=>` Most of the trends in chemical properties of elements, such as diagonal relationships, inert pair effect, effects of lanthanoid contraction etc. will be dealt with along the discussion of each group in later units.

`=>` In this section we shall study the periodicity of the valence state shown by elements and the anomalous properties of the second period elements (from lithium to fluorine).

`=>` The valence is the most characteristic property of the elements and can be understood in terms of their electronic configurations.

`=>` The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and/or equal to eight minus the number of outermost electrons as shown in fig.

`=>` Nowadays the term oxidation state is frequently used for valence.

`=>` Consider the two oxygen containing compounds : `OF_2` and `Na_2 O`.

● The order of electronegativity of the three elements involved in these compounds is `F > O > Na`.

● Each of the atoms of fluorine, with outer electronic configuration `2s^2 2p^5`, shares one electron with oxygen in the `OF_2` molecule.

● Being highest electronegative element, fluorine is given oxidation state `–1`.

● Since there are two fluorine atoms in this molecule, oxygen with outer electronic configuration `2s^2 2p^4` shares two electrons with fluorine atoms and thereby exhibits oxidation state `+2`.

● In `Na_2O`, oxygen being more electronegative accepts two electrons, one from each of the two sodium atoms and, thus, shows oxidation state `–2`.

● On the other hand sodium with electronic configuration `3s^1` loses one electron to oxygen and is given oxidation state `+1`.

● Thus, the oxidation state of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.


`=>` Some periodic trends observed in the valence of elements (hydrides and oxides) are shown in Table 3.9.

`=>` There are many elements which exhibit variable valence.

● This is particularly characteristic of transition elements and actinoids.

`=>` The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other members of their respective group.

`text(Example :)` ● Lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals, form compounds with pronounced covalent character; the other members of these groups predominantly form ionic compounds.

`text(Diagonal Relationship :)` The behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.

`=>` The anomalous behaviour is attributed to :

(i) small size

(ii) large charge/ radius ratio

(iii) high electronegativity of the elements.

(iv) The first member of group has only four valence orbitals (`2s` and `2p`) available for bonding, whereas the second member of the groups have nine valence orbitals (`3s`, `3p`, `3d`).

● As a consequence of this, the maximum covalency of the first member of each group is `4` (e.g., boron can only form `[BF_4]^-`, whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons e.g., aluminium forms `[AlF_6]^(3-)`.

(v) Furthermore, the first member of `p`-block elements displays greater ability to form `pπ – pπ` multiple bonds to itself (e.g., `C = C`, `C ≡ C`, `N = N`, `N ≡ Ν`) and to other second period elements (e.g., `C = O`, `C = N`, `C ≡ N`, `N = O`) compared to subsequent members of the same group.

`=>` The ionization enthalpy of the extreme left element in a period is the least and the electron gain enthalpy of the element on the extreme right is the highest negative. This results into high chemical reactivity at the two extremes and the lowest in the centre.

`=>` The maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation and at the extreme right (among halogens) shown by the gain of an electron forming an anion.

● This property can be related with the reducing and oxidizing behaviour of the elements.

● It can be directly related to the metallic and non-metallic character of elements.

● Thus, the metallic character of an element, which is highest at the extremely left decreases and the non-metallic character increases while moving from left to right across the period.

`=>` The chemical reactivity of an element can be best shown by its reactions with oxygen and halogens.

● Elements on two extremes of a period easily combine with oxygen to form oxides.

● The normal oxide formed by the element on extreme left is the most basic (e.g., `Na_2O`), whereas that formed by the element on extreme right is the most acidic (e.g., `Cl_2O_7`).

● Oxides of elements in the centre are amphoteric (e.g., `Al_2O_3`, `As_2O_3`) or neutral (e.g., `CO`, `NO`, `N_2O`).

● Amphoteric oxides behave as acidic with bases and as basic with acids, whereas neutral oxides have no acidic or basic properties.

`=>` Among transition metals (3d series), the change in atomic radii is much smaller as compared to those of representative elements across the period.

`=>` The change in atomic radii is still smaller among inner-transition metals (4f series).

● The ionization enthalpies are intermediate between those of s- and p-blocks.

● As a consequence, they are less electropositive than group 1 and 2 metals.

`=>` In a group, the increase in atomic and ionic radii with increase in atomic number generally results in a gradual decrease in ionization enthalpies and a regular decrease (with exception in some third period elements) in electron gain enthalpies in the case of main group elements.

● Thus, the metallic character increases down the group and non-metallic character decreases. This trend can be related with their reducing and oxidizing property.

● In the case of transition elements, however, a reverse trend is observed. This can be explained in terms of atomic size and ionization enthalpy.