Classification of Elements and Periodicity in PropertiesMind Map

Periodic classification was needed because the number of known elements increased rapidly and memorizing each element separately became impossible. Scientists tried to arrange elements so that similarities in physical and chemical properties became visible. Dobereiner grouped elements into triads, where the atomic mass of the middle element was approximately the average of the other two. Newlands arranged elements by increasing atomic mass and observed repetition after every eighth element, called the law of octaves. Mendeleev gave the first highly successful periodic table based mainly on atomic mass and chemical properties. He left gaps for undiscovered elements and predicted their properties, but his table could not properly explain isotopes, anomalous pairs and the position of hydrogen.

The modern periodic table is based on the modern periodic law: physical and chemical properties of elements are periodic functions of their atomic numbers. Since atomic number determines electronic configuration, the table becomes a systematic map of electron arrangement and properties. The long form periodic table has 7 periods and 18 groups. Periods are horizontal rows and generally indicate the highest occupied shell, while groups are vertical columns containing elements with similar valence-shell configurations. Elements are also positioned into s, p, d and f blocks. For elements with atomic number above 100, IUPAC temporary nomenclature uses numerical roots such as nil, un, bi and tri. NEET questions commonly ask group, period, block and valence configuration from atomic number.

Electronic configuration describes how electrons are distributed among shells, subshells and orbitals. It is the bridge between atomic structure and the periodic table because the position and properties of an element depend mainly on its valence-shell configuration. Electrons fill orbitals according to the Aufbau principle, which follows increasing energy based on the n + l rule. Pauli’s exclusion principle states that an orbital can hold a maximum of two electrons with opposite spins. Hund’s rule says that electrons occupy degenerate orbitals singly with parallel spins before pairing. Valence electrons decide group, block, valency, bonding behavior and chemical reactivity. Exceptions such as chromium and copper occur because half-filled and fully filled subshells give extra stability.

Elements are classified into s, p, d and f blocks according to the subshell into which the differentiating electron enters. This block classification connects electronic configuration with chemical behavior. s-block elements have their last electron in an s orbital and include Groups 1 and 2, except helium’s special placement. p-block elements have their last electron in a p orbital and include Groups 13 to 18; they contain metals, non-metals and metalloids. d-block elements are transition elements where the penultimate d subshell is being filled, commonly showing variable oxidation states and colored compounds. f-block elements are inner transition elements involving 4f and 5f filling, including lanthanoids and actinoids. NEET often asks block, general configuration and characteristic trends.

Periodic trends are regular variations in properties caused by changing effective nuclear charge, number of shells and shielding. Across a period, electrons enter the same shell while nuclear charge increases, so effective nuclear charge increases and atomic radius generally decreases. Down a group, new shells are added and shielding increases, so size increases. Ionization enthalpy generally increases across a period and decreases down a group, with important exceptions such as Be-B and N-O. Electron gain enthalpy generally becomes more negative across a period, but chlorine is more negative than fluorine due to lower electron-electron repulsion. Electronegativity increases across and decreases down. Valency, metallic character, non-metallic character and oxidizing or reducing nature follow from electron loss or gain tendencies.