Classification of Elements and Periodicity in Properties: Periodic Table

Introduction

• The periodic table arranges all known elements based on their properties, grouping those with similar attributes together.

Development of the Periodic Table

• Lavoisier Classification: Elements classified as metals and non-metals.

  • Metals lose electrons (e.g., Na → Na⁺ + e⁻), non-metals gain electrons (e.g., F + e⁻ → F⁻).
  • Limitations: Increase in elements made classification insufficient. Metalloids posed classification challenges.

• Dobereiner's Triad Rule (1817)

  • Groups of three elements with similar properties.
  • Middle element's atomic weight averages the other two.
  • Example Triads: (Li, Na, K), (Cl, Br, I).
  • Limitation: Unable to arrange all known elements as triads.

• Newland's Octave Rule (1865)

  • Elements arranged by increasing atomic mass; every 8th element shows similar properties.
  • Rule held until elements with d-block differences emerged.

• Lothar Meyer's Curve (1869)

  • Curve between atomic weight and volume, showing periodicity of properties such as electropositivity and electronegativity.

• Mendeleev's Periodic Table

  • Based on atomic weight.
  • 63 known elements, organized by properties.
  • Elements arranged in periods (horizontal) and groups (vertical).
  • Proposed positions for undiscovered elements based on properties.
  • Limitations: Misplacement of elements with similar properties, unclear placement of isotopes, and anomalies with hydrogen.

Modern Periodic Table

• Moseley’s Contribution

  • Introduced atomic number as the basis for arrangement.
  • Modern periodic law states that properties are periodic functions of atomic number.

• Present Form

  • Organized by Bohr-Bury electronic configuration.
  • 7 periods and 18 groups.
  • Similar group elements share outer electron configurations.

Periodic Trends

• Atomic Radius: Average distance from nucleus to valence shell.

  • Decreases across a period due to increased nuclear charge.
  • Increases down a group due to added shells.

• Ionisation Energy: Energy to remove an electron.

  • Increases across a period due to increased nuclear attraction.
  • Decreases down a group as distance from nucleus increases.

• Electron Affinity: Energy change when an electron is added.

  • First addition typically exothermic; subsequent can be endothermic.
  • Varies based on atomic size and configuration stability.

• Electronegativity: Tendency of an atom to attract shared electrons.

  • Increases across a period, decreases down a group.
  • Affects bond polarity and molecular characteristics.

Key Concepts

• Magic Numbers: 2, 8, 18, 32 - intervals at which properties repeat.
• Shielding Effect: Inner electrons reduce nucleus' pull on outer electrons.
• Lanthanide Contraction: Poor shielding by f-electrons leads to decreasing size across the series.

Application of Periodic Trends

• Metallic and Non-metallic Nature: Metals less electronegative than non-metals.
• Reactivity: Increases with lower ionization energy for metals.
• Stability Predictions: Energy differences reveal stable states.

These notes summarize the classification and periodicity of elements in the periodic table, focusing on historical development, modern updates, periodic trends, and their applications in predicting element behavior and properties.