PRELIMS

Atomic Structure

Basic Definitions

• Atom: Smallest particle of an element that takes part in chemical combination; does not occur free in nature. (Smallest/lightest: Hydrogen).
• Molecule: Smallest particle of an element or compound capable of stable, independent existence (e.g., ).
• Mole: A collection of particles.
• Avogadro's Number: .

Mass Concepts

• Atomic Mass: Ratio of the mass of one atom of an element to th of the mass of a Carbon-12 atom.
  • Actual Mass formula: Atomic mass (amu) .
• Molecular Mass: Indicates how many times a molecule is heavier than th the mass of Carbon-12.

Fundamental Particles

Electron

• Discoverer: J.J. Thomson | Named by: Stoney
• Charge: -1 (relative), C (absolute)
• Mass: g

Proton

• Discoverer: Goldstein | Named by: Rutherford
• Charge: +1 (relative), C (absolute)
• Mass: g

Neutron

• Discoverer: James Chadwick
• Charge: 0 (neutral)
• Mass: g

• Atomic Number (Z): Number of protons ( ) or electrons ( ) in a neutral atom. .
• Mass Number (A): Total number of protons and neutrons. .

Isotopes, Isobars, Isotones, and Isoelectronic

• Isotopes: Same atomic number (Z), different mass number (A). (e.g., ).
• Isobars: Same mass number (A), different atomic numbers (Z). (e.g., ).
• Isotones: Different elements with the same number of neutrons. (e.g., - all have 8 neutrons).
• Isoelectronic: Atoms/ions/molecules with the same number of electrons. (e.g., ).

Quantum Numbers (Basics)

Four numbers that describe the position and energy of an electron in an atom.

• Principal Quantum Number (n): Shell number (K, L, M, N... = 1, 2, 3, 4...); determines size/energy of orbital.
• Azimuthal/Angular Momentum Quantum Number (l): Sub-shell (s, p, d, f = 0, 1, 2, 3); determines shape of orbital. Range: to .
• Magnetic Quantum Number (m): Orientation of orbital in space. Range: to .
• Spin Quantum Number (s): Spin of electron: or .
• Maximum electrons in a shell: (K=2, L=8, M=18, N=32).

Electron Configuration Rules

• Aufbau Principle: Electrons fill orbitals in order of increasing energy (lower energy orbitals filled first). Order:
• Pauli's Exclusion Principle: No two electrons in an atom can have the same set of all four quantum numbers; an orbital can hold a maximum of 2 electrons with opposite spins.
• Hund's Rule of Maximum Multiplicity: Electrons occupy degenerate (same-energy) orbitals singly first, with parallel spins, before pairing up.

Historical Models of the Atom

1. Thomson's Model (Plum-Pudding/Watermelon): Atom is a sphere (radius ) of uniform positive charge with embedded electrons.
2. Rutherford's Nuclear Model: Based on scattering experiments.
   • Atom has a heavy positive nucleus (radius ) containing nucleons (protons + neutrons).
   • Atom volume is times the nucleus volume.
   • Electrons revolve in closed orbits like planets around the sun (Planetary electrons).
3. Bohr's Model: Concept of quantization of energy.
4. Modern Theories:
   • Planck's Quantum Theory: Light/Energy in photons/quanta.
   • de-Broglie Equation: Dual nature of electrons (wave and particle).
   • Heisenberg's Uncertainty Principle: Exact position and momentum cannot be determined simultaneously.
   • Schrödinger's Wave Equation: Wave nature of electron.

Key Phenomena

• Spectral Lines: Splitting of light through a prism.
• Zeeman's Effect: Splitting of spectral lines in a magnetic field.
• Stark's Effect: Splitting of spectral lines in an electric field.

UPSC Relevance

Atomic structure fundamentals are essential for understanding chemistry and advanced S&T (like nanotechnology).

• Models: Historical progression of atomic models (Thomson to Schrödinger).
• Definitions: Distinguishing between isotopes, isobars, and especially isotones/isoelectronic ions is a common Prelims check.
• Radioactivity Link: Nucleons (protons/neutrons) and their role in nuclear stability.