Revision Notes: Classification of Elements & Periodicity in Properties | Chemistry Class 11 - NEET PDF Download

Introduction

The periodic table organizes elements based on their properties, evolving from early classifications to a systematic arrangement reflecting atomic structure and periodicity in physical and chemical properties.

Genesis of Periodic Classification

• Johann Dobereiner (1829): Proposed Triads—groups of three elements with the middle element’s atomic weight averaging the other two and properties intermediate (e.g., Li, Na, K).Dobereiner’s Triads
• A.E.B. de Chancourtois (1862): Arranged elements by increasing atomic weights on a cylindrical table, showing periodic recurrence.
• John Newlands (1865): Law of Octaves—every eighth element repeats properties, valid up to calcium (e.g., Li to F, Na to Cl).Newlands’ Octaves
• Lothar Meyer (1869): Plotted physical properties (e.g., atomic volume) vs. atomic weight, noting periodic patterns.
• Dmitri Mendeleev (1869): Formulated Periodic Law—“Properties of elements are periodic functions of their atomic weights.” Arranged elements in a table, prioritizing similar properties over strict atomic weight order, predicting undiscovered elements (e.g., Eka-aluminium, Eka-silicon).

Mendeleev’s Periodic Table

• Arranged elements by increasing atomic weights, grouping those with similar properties (e.g., Iodine with halogens despite lower atomic weight than Tellurium).
• Predicted properties of undiscovered elements (e.g., gallium, germanium), validated later:Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) andEka-silicon (Germanium)

Modern Periodic Law and Table

• Henry Moseley (1913): Showed atomic number (Z) as the fundamental property via X-ray spectra, leading to Modern Periodic Law—“Properties are periodic functions of atomic numbers.”
• Structure: 7 periods (rows), 18 groups (columns); periods reflect principal quantum number (n), groups reflect similar valence electron configurations.
• Periods: 1 (2 elements), 2-3 (8 each), 4-5 (18 each), 6-7 (32 each, 7th incomplete).
• Groups: Numbered 1-18 (IUPAC); lanthanoids and actinoids placed separately.

Nomenclature of Elements (Z > 100)

• IUPAC systematic naming uses numerical roots until discovery is confirmed (e.g., Z=120: unbinilium, Ubn).
• Notation:
• Examples: Z=106 (Seaborgium, Sg), Z=114 (Flerovium, Fl).

Electronic Configurations and Periodic Table

• Periods: Reflect filling of shells (n); e.g., Period 1 (1s, 2 elements), Period 4 (4s, 3d, 4p, 18 elements).
• Groups: Similar valence configurations (e.g., Group 1: ns¹).
• Blocks:
  • s-block: Groups 1-2, filling s-orbitals.
  • p-block: Groups 13-18, filling p-orbitals.
  • d-block: Groups 3-12, filling d-orbitals.
  • f-block: Lanthanoids/Actinoids, filling f-orbitals.
• Exceptions: H (1s¹, unique), He (1s², with noble gases).

Types of Elements

• Metals: >78% of elements, left side, high melting/boiling points, conductors, malleable, ductile.
• Non-metals: Top-right, low melting/boiling points, poor conductors, brittle.
• Metalloids: Borderline (e.g., Si, Ge), mixed properties.
• Trends: Metallic character increases down groups, decreases across periods.

Periodic Trends in Physical Properties

Atomic Radius:

• Decreases across period (increased nuclear charge); e.g., Li (152 pm) to F (64 pm).
• Increases down group (more shells); e.g., Li (152 pm) to Cs (262 pm).

Ionic Radius: 

• Cations smaller (e.g., Na⁺ 95 pm vs. Na 186 pm), anions larger (e.g., F⁻ 136 pm vs. F 64 pm); 
• isoelectronic species vary by nuclear charge (e.g., Mg²⁺ < Na⁺ < F⁻).

Ionization Enthalpy (ΔᵢH):

• Increases across period (e.g., Na 496 to Si 786 kJ mol⁻¹), decreases down group (e.g., Li to Cs).
• Exceptions: B < Be (s vs. p), O < N (electron repulsion).

Electron Gain Enthalpy (Δₑ₉H):

• More negative across period (e.g., F -328 kJ mol⁻¹), less negative down group (e.g., Cl -349 to I -295 kJ mol⁻¹).
• Exceptions: O, F less negative than S, Cl (repulsion in n=2).

Electronegativity:

• Increases across period (e.g., Li 1.0 to F 4.0), decreases down group (e.g., F 4.0 to At 2.2).

Periodic Trends in Chemical Properties

• Valence/Oxidation State: Equals valence electrons or 8 minus valence electrons (e.g., Group 1: 1, Group 17: 1 or 7); e.g., OF₂ (O +2), Na₂O (Na +1, O -2).
• Anomalous Properties of Second Period: Li, Be, B to F differ due to small size, high electronegativity, limited orbitals (max covalency 4); e.g., [BF₄]⁻ vs. [AlF₆]³⁻.
• Chemical Reactivity: High at extremes (Group 1 loses e⁻, Group 17 gains e⁻); oxides basic (left, e.g., Na₂O) to acidic (right, e.g., Cl₂O₇), amphoteric/neutral in center (e.g., Al₂O₃, CO).