Blocks in the Periodic Table

Elements in the periodic table can also be classified into four blocks. This classification Is based upon the valence orbital of the element involved in chemical bonding. According to this Classification, elements of IA and IIA subgroups are called s-block elements because their Valence electrons are available in s orbital. The elements of IIIA to VIIIA subgroups (except He) Are known as p-block elements as their valence electrons are present in p orbital. Similarly in Transition elements, electrons in d-orbital are responsible for their valency hence they are called d-block elements. For Lanthanides and Actinides valence electrons are present in f- orbital hence these elements are called f-block elements. This classification is quite useful in understanding the chemistry of elements and predicting their properties especially the concept of valency or oxidation state.

Metals, Non-metals and Metalloids.

Another basis for classifying the elements in the periodic table is their metallic character. Generally, the elements on the left hand side, in the centre and at the bottom of the periodic table Are metals, while the non-metals are in the upper right corner of the table. Some elements, Especially lower members of groups, IVA, VA and VIA (as shown in Table 1.1) have properties of Both metals as well as non-metals.

 These elements are called semi-metals or metalloids. In thePeriodic table elements of groups IVA to VIIIA, at the top right hand corner above the stepped Line, are non-metals. The elements just under the “steps’ such as Si, As, and Te are the metalloids. All the remaining elements, except hydrogen, are metals.

PERIODIC TRENDS IN PHYSICAL PROPERTIES

As you have studied so far that in modern periodic table the elements are arranged in Ascending order of their atomic numbers and their classification in groups and periods is based o The similarity in their properties. Yet, due to the gradual increase in the number of protons in the Nucleus and electrons in outer shells the physical and chemical properties of the elements Steadily vary within a group or a period. Here, we study some trends in physical properties.

ATOMIC SIZE

Atomic Radius:

Atoms are so small that it is impossible to see an atom even with a powerful optical Microscope. The size of a single atom therefore cannot be directly measured. However, Techniques have been developed which can measure the distance between the centres of two Bonded atoms of any element. Half of this distance is considered to be the radius of the atom. In the periodic table, the atomic radius increases from top to bottom within a group due to Increase in atomic number. 

This is because of the addition of an extra shell of electrons in each Period. In a period, however, as the atomic number increases from left to right, the atomic radius Decreases. This gradual decrease in the radius is due to increase in the positive charge in the Nucleus. As the positive nuclear charge increases, the negatively charged electrons in the shells Are pulled closer to the nucleus. Thus, the size of the outermost shell becomes gradually smaller. 

This effect is quite remarkable in the elements of longer periods in which “d” and “f” subshells Are involved. For example, the gradual decrease in the size of Lanthanides is significant and Called Lanthanide Contraction.

Ionic Radius:

When a neutral atom loses one or more electrons, it becomes a positive ion. The size of The atom is decreased in this process because of the two reasons. First the removal of one or more Electrons from a neutral atom usually results in The loss of the outermost shell and second, the Removal of electrons causes an imbalance in Proton-electron ratio. Due to the greater Attraction of the nuclear charge, the remaining Electrons of the ion are drawn closer to the Nucleus. Thus, a positive ion is always smaller Than the neutral atom from which it is derived. The radius of Na is 186 pm and the radius of Na Is 102 pm. On the contrary, a negative ion is Always bigger than its parent atom. 

The reason Is that addition of one or more electrons in the Shell of a neutral atom enhances repulsion Between the electrons causing expansion of the Shell. Thus, the radius of fluorine atom is 72 pm And that of the fluoride ion (F) is 133 pm. In a group of the periodic table, similar Charged ions increase in size from top to Bottom. 

Whereas within a period, isoelectronic Positive ions show a decrease in ionic radius From left to right, because of the increasing Nuclear charge. The same trend is observed for The isoelectronic negative ions of a period; ionic Size decrease from left to right. The variations In atomic and ionic radii of alkali metals and Halogens are shown in Fig 1.1 and Fig. 1.2.

Ionization Energy

The ionization energy of an element is the minimum quantity of energy which is required To remove an electron from the outermost shell of its isolated gaseous atom in its ground state. Elements with greater number of electrons have more than one values of ionization.

Variation Within a Group: 

The factors upon which the ionization energy of an atom mainly depends are magnitude Of nuclear charge, size of the atom, and the “shielding effect”. The shielding effect is actually the Repulsion due to electrons in between the nucleus And the outermost shell. This effect increases, as the Size of the atom increases due to addition of an extra Shell successively in each period hence more Number of electrons shields the nucleus. Going Down in a group, the nuclear charge increases but as The size of the atom and the number of electrons Causing the shielding effect also increases therefore Ionization energy decreases from top to bottom. 

That Is why in alkali metals, for example, it is easier to Remove an electron from caesium atom than from Lithium atom. The change in ionization energies of IA elements is shown in Fig. 1.3.

Variation Across a Period:

Generally, smaller the atom with greater Nuclear charge, more strongly the electrons are Bound to the nucleus and hence higher the i Ionization energy of the atom. By moving from left To right in a period, the outer shell remains the Same, while the nuclear charge increases Effectively that makes the removal of an electron Difficult and hence the value of ionization energy Increases.

 Although, the number of electrons also Increases in this case but the shielding is not very Effective within the same shell. The trend of Ionization energies of short periods is shown in Fig.1.4. The figure also reveals that inert gases have The highest values of ionization energy because due To complete outermost shell in them, the removal of Electron is extremely difficult.