Nonmetals usually react by gaining electrons.

The non-metallic character increases while moving from left to right across the period.

Hence, the reactivity increases down the group. The metallic character increases down the group. Thus it is easier for those electrons to come off. Consequently the dominant factor is that we have more and more energy levels and the electrons are further and further away from the nucleus.

As you go down the periodic table, the nuclear charge increases but so does the number of shielding electrons.

The metallic character of an element, which is highest at the extremely left decreases from left to right across the period as the electrons are more tightly bound due to increased nuclear charge.

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

As a consequence, they are less electropositive and thus less reactive than group 1 and 2 metals. The ionization enthalpies are intermediate between those of s- and p-blocks.

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

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

This is because the elements present at left extreme of a period have the lowest ionization enthalpy and can lose an electron to form a cation and that present at the right extreme have the highest electron affinity and can gain an electron to form an anion.

Chemical reactivity is highest at the two extremes of a period and is lowest at the center.

Non availability of d-orbitals for bonding.

Element of second period are known as bridge elements. These pairs of element have almost identical ionic radii and polarizing power (i.e. Anomalous properties of second period elements – diagonal relationshipĬertain elements in the 2 nd period show similarity with their diagonal elements in 3 rd period. The stability of the +1 oxidation state increases in the following sequence:Ģ.

As an example in group 13 the +1 oxidation state of Tl is the most stable and Tl3+ compounds are comparatively rare.

The term inert pair effect is often used in relation to the increasing stability of oxidation states that are 2 less than the group valency for the heavier elements of groups 13, 14, 15 and 16.

The inert pair effect is the tendency of the electrons in the outermost atomic s-orbitalto remain unionized or unshared in compounds of post-transition metals.

In general, the stability of the higher oxidation states is in the order 3d << 4d << 5d Oxidation state = (n − 1) d e −(unpaired) +ns e −.

Higher p-block elements exhibit more stability for lower oxidation state due to inert pair effect.

In case of alkali and alkaline earth metals it is fixed i.e., + 1 and + 2 respectively.

Its value can be zero, positive, negative or fractional.

The oxidation state of O is -2 and that of Na is +1 in Na₂F. the oxidation state of F is -1 and O is +2 in OF₂.

For example: in the compounds OF₂ and Na 2O the electronegativity order for the elements is F > O > Na.It is defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms.Nowadays the term oxidation state is frequently used for valence.Many elements, particularly transition elements and actinoids, exhibit variable valency as shown in the table below.

Going across a period from left to right the number of valence electrons increases.The valency of all the elements in a group is same as the electrons in outer shell for elements of a group is same as shown below.The valence of 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.Thevalence (or valency) of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules.Periodicity in valency or oxidation states:

Reactivity trend across periodic table