Revision Notes: Classification of Elements & Periodicity in Properties

Some Important Points and Terms of the Chapter

5. Modern periodic law: Properties of elements are the periodic function of their atomic number.
6. Modern Periodic Table: This table was prepared was Bohr and is based upon the electronic configuration of elements. The table consists of 18 vertical columns called groups. Elements having similar outer electronic configurations in their atoms are arranged in vertical columns, referred to as groups. According to the recommendation of International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 and the table consists of 7 horizontal rows called periods. The first period contains 2 elements. The subsequent periods consists of 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and like the sixth period would have a theoretical maximum (on the basis of quantum numbers) of 32 elements. In this form of the Periodic Table, 14 elements of both sixth and seventh periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom

7. Notation for IUPAC Nomenclature of Elements With Z > 100
8. We can classify the elements into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbital that are being filled with electrons.
9. s-Block Elements : The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns¹ and ns² outermost electronic configuration belong to the s-Block Elements.
10. p-Block Elements The p-Block Elements comprise those belonging to Group 13 to 18 and these together with the s-Block Elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns²np¹ to ns²np⁶ in each period.
11. d-Block Elements These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n-1)d^1-10 ns^0-2.
12. f-Block Elements The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) - Lu(Z = 71) and Actinoids, Th(Z = 90) - Lr (Z = 103) are characterised by the outer electronic configuration (n-2)f^1-14 (n-1)d^0-1 ns². The last electron added to each element is filled in f- orbital. These two series of elements are hence called the Inner-Transition Elements (f-Block Elements).
13. Variation in Atomic Radius in Period : The atomic size generally decreases across a period It is because within the period the outer electrons are in the same valence shell and the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.
    

14. Variation in Atomic Radius in Group: Within a family or vertical column of the periodic table, the atomic radius increases regularly with atomic number as), as we descend the groups, the principal quantum number (n) increases and the valence electrons are farther from the nucleus. This happens because the inner energy levels are filled with electrons, which serve to shield the outer electrons from the pull of the nucleus. Consequently the size of the atom increases as reflected in the atomic radii.
15. The atomic radii of noble gases are not considered here. Being monatomic, their (non-bonded radii) values are very large. In fact radii of noble gases should be compared not with the covalent radii but with the van der Waals radii of other elements.
16. A cation is smaller than its parent atom because it has fewer electrons while its nuclear charge remains the same. The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge. For example, the ionic radius of fluoride ion (F^-) is 136 pm whereas the atomic radius of fluorine is only 64 pm. On the other hand, the atomic radius of sodium is 186 pm compared to the ionic radius of 95 pm for Na⁺.
17. Isoelectronic species : Atoms and ions which contain the same number of electrons.. For example, O^2-, F^-, Na⁺ and Mg²⁺ have the same number of electrons (10). Their radii would be different because of their different nuclear charges. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative charge will have the larger radius. In this case, the net repulsion of the electrons will outweigh the nuclear charge and the ion will expand in size.
18. Ionization Enthalpy: It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state. In other words, the first ionization enthalpy for an element X is the enthalpy change (ΔH) for the reaction depicted in equation. X (g) → X⁺(g) + e^-. The ionization enthalpy is expressed in units of kJ mol^-1. We can define the second ionization enthalpy as the energy required to remove the second most loosely bound electron; it is the energy required to carry out the reaction shown in equation X⁺(g) → X²⁺(g) + e^-. Energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive. The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom.