Chapter 3: Classification of Elements and Periodicity in Properties | Summary

In 1913, Henry Moseley showed through X-ray spectra that a plot of the square root of X-ray frequency against atomic number gave a straight line, demonstrating that atomic number is more fundamental than atomic mass. Mendeleev's law was modified to the Modern Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic numbers.

The modern long form of the periodic table has 7 periods and 18 groups. The period number corresponds to the highest principal quantum number of elements in that period. Periods contain 2, 8, 8, 18, 18, and 32 elements respectively. Groups are numbered 1 to 18 as per IUPAC recommendations. Elements with similar outer electronic configurations appear in the same vertical column. The 14 lanthanoid and actinoid elements are placed separately at the bottom.

New elements with very high atomic numbers are extremely unstable and exist in only minute quantities. To avoid disputes over naming rights, IUPAC recommends a systematic nomenclature derived directly from the atomic number using numerical roots: nil(0), un(1), bi(2), tri(3), quad(4), pent(5), hex(6), sept(7), oct(8), enn(9). The roots are combined in order of digits and 'ium' is added at the end.

For example, element 120 would be named unbinilium with symbol Ubn. Once officially discovered and verified, a permanent name reflecting the country of discovery or a notable scientist is given by IUPAC vote. Elements up to atomic number 118 have been discovered and officially named, for example Oganesson (118), Tennessine (117), and Seaborgium (106).

An element's position in the periodic table directly reflects the quantum numbers of its last filled orbital. Each period corresponds to filling of a new principal energy level. The first period has 2 elements (1s filled), the second and third have 8 each, and the fourth and fifth have 18 each due to filling of d orbitals. The sixth period contains 32 elements with filling of 4f orbitals giving the lanthanoid series.

Elements in the same group share the same valence shell electronic configuration, explaining their similar chemical behaviour. The 3d transition series starts at scandium (Z=21) and the 4f lanthanoid series starts at cerium (Z=58). The 5f actinoid series starts at thorium (Z=90). The fourth period contains 18 elements because when n=4, orbitals 4s, 3d, and 4p can together accommodate 18 electrons.

Elements are classified into four blocks. The s-block (Groups 1 and 2) consists of alkali and alkaline earth metals with ns1 or ns2 configurations; they are highly reactive and never found free in nature. The p-block (Groups 13-18) includes representative elements with ns2np1 to ns2np6 configurations; it includes non-metals, metalloids, and noble gases. Halogen (Group 17) and chalcogen (Group 16) elements have highly negative electron gain enthalpies.

The d-block (Groups 3-12) or transition elements have general configuration (n-1)d1-10ns0-2; they form coloured ions, show variable oxidation states, paramagnetism, and catalytic activity. The f-block elements (lanthanoids Ce-Lu and actinoids Th-Lr) are inner transition elements with (n-2)f1-14(n-1)d0-1ns2 configuration; they are all metals. Helium, though an s-block element, is placed in the p-block with noble gases due to its completely filled valence shell. Hydrogen is placed separately due to its dual character resembling both alkali metals and halogens.

Metals constitute more than 78% of known elements and appear on the left side of the periodic table. They are usually solids at room temperature (mercury is an exception), have high melting and boiling points, are good conductors of heat and electricity, and are malleable and ductile. Non-metals are located at the top right of the periodic table, are usually solids or gases with low melting points, and are poor conductors.

The change from metallic to non-metallic character across a period is gradual, represented by a thick zig-zag line in the periodic table. Elements bordering this line, such as silicon, germanium, arsenic, antimony, and tellurium, show characteristics of both metals and non-metals and are called metalloids or semi-metals. Metallic character increases down a group and decreases across a period from left to right.

Atomic radius decreases across a period because the effective nuclear charge increases while electrons are added to the same shell, pulling them closer to the nucleus. Atomic radius increases down a group because a new principal energy level is added, and inner electrons shield outer electrons from the nucleus. Ionic radius follows similar trends: cations are smaller than parent atoms (fewer electrons, same nuclear charge) while anions are larger (extra electrons increase repulsion).

Ionization enthalpy generally increases across a period (increasing nuclear charge, decreasing radius) and decreases down a group (increasing shielding). Exceptions include the lower ionization enthalpy of boron compared to beryllium (2p electron more shielded than 2s) and oxygen compared to nitrogen (extra electron in one 2p orbital causes repulsion). Electron gain enthalpy becomes more negative across a period and less negative down a group, with exceptions for O and F due to small 2p orbital repulsion. Electronegativity follows the same trend as ionization enthalpy, with fluorine being assigned 4.0 on the Pauling scale.

The valence of representative elements equals either the number of outermost electrons or eight minus that number. Transition elements and actinoids show variable valence. The first member of each group in s- and p-blocks (e.g., Li, Be, B to F) shows anomalous properties due to small size, high charge-to-radius ratio, high electronegativity, and availability of only four valence orbitals (2s and 2p), restricting maximum covalency to 4.

A diagonal relationship exists in the periodic table: lithium resembles magnesium, and beryllium resembles aluminium in their chemical properties. First period p-block elements also show unique ability to form p-pi to p-pi multiple bonds. Across a period, metallic character decreases and non-metallic character increases. Elements on the left extreme of a period form basic oxides (e.g., Na2O) while those on the right form acidic oxides (e.g., Cl2O7); central elements form amphoteric or neutral oxides.

The periodic table is the most powerful organising principle in chemistry, showing that elements are not random but display systematic trends in physical and chemical properties. From Dobereiner's triads to Mendeleev's predictions and Moseley's modern periodic law, the development of the periodic table represents the cumulative effort of scientists across generations. The arrangement of 118 known elements into s, p, d, and f blocks based on electronic configuration provides a rational framework for understanding and predicting chemical behaviour.

For CBSE board examinations, this chapter is highly important as it forms the foundation for all subsequent topics in inorganic chemistry. Students must master periodic trends in atomic radius, ionization enthalpy, electron gain enthalpy, and electronegativity, as well as block classification, IUPAC nomenclature for elements above Z=100, and the anomalous behaviour of second period elements. Understanding these trends enables prediction of chemical reactivity, bond character, and compound formation, making this chapter essential for both theory and application-based questions.