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1. Atomic Structure

Definition: Atomic structure refers to the composition and arrangement of subatomic particles (protons, neutrons, and electrons) within an atom.

Subatomic Particles:

  • Protons (p⁺): Positively charged particles found in the nucleus.
  • Neutrons (n⁰): Neutral particles also located in the nucleus.
  • Electrons (e⁻): Negatively charged particles revolving around the nucleus in defined orbits or shells.

Atomic Models:

  • Dalton’s Atomic Theory (1803): Atoms are indivisible and indestructible.
  • Thomson’s Model (1897): Plum pudding model; atoms consist of a positively charged sphere with embedded electrons.
  • Rutherford’s Model (1911): Nucleus at the center with electrons revolving around it.
  • Bohr’s Model (1913): Electrons revolve in discrete orbits with fixed energy levels.
  • Quantum Mechanical Model: Proposed by Schrödinger; uses wave functions to describe the probability of finding an electron in a particular region around the nucleus.

Quantum Numbers:

  1. Principal Quantum Number (n): Describes the size and energy of the orbit.
  2. Azimuthal Quantum Number (l): Defines the shape of the orbital (s, p, d, f).
  3. Magnetic Quantum Number (m): Specifies the orientation of the orbital.
  4. Spin Quantum Number (s): Represents the spin of the electron (±1/2).

Electronic Configuration:

  • Aufbau Principle: Electrons fill orbitals in the order of increasing energy (1s < 2s < 2p < 3s …).
  • Pauli’s Exclusion Principle: No two electrons can have all four quantum numbers identical.
  • Hund’s Rule: Every orbital in a subshell gets one electron before any gets two.

2. Periodic Properties of Elements

Periodic Table:

  • Modern Periodic Law: The properties of elements are a periodic function of their atomic numbers.
  • Groups and Periods: 18 vertical columns (groups) and 7 horizontal rows (periods).

Periodic Trends:

  1. Atomic Radius:
    • Across a period: Decreases due to increased nuclear charge pulling electrons closer.
    • Down a group: Increases due to the addition of shells.
  2. Ionization Energy (IE): Energy required to remove an electron from a neutral atom.
    • Across a period: Increases as nuclear attraction strengthens.
    • Down a group: Decreases as outer electrons are further from the nucleus.
  3. Electron Affinity (EA): Energy released when an electron is added to a neutral atom.
    • Across a period: Generally increases.
    • Down a group: Generally decreases.
  4. Electronegativity: Tendency of an atom to attract a bonding pair of electrons.
    • Across a period: Increases.
    • Down a group: Decreases.
  5. Metallic and Non-metallic Character:
    • Metallic character: Increases down a group, decreases across a period.
    • Non-metallic character: Increases across a period, decreases down a group.
  6. Valency: Number of electrons an atom gains, loses, or shares to achieve a stable configuration.

Anomalies:

  • Transition elements show variable valency due to the involvement of d-orbitals.
  • Lanthanide and actinide contraction affect atomic and ionic sizes.

3. Significance of Periodic Properties

  • Helps predict chemical behavior of elements.
  • Aids in understanding trends in reactivity, stability, and bonding.
  • Crucial for the development of new materials and compounds.

4. Applications in APSC Preparation

  • Questions are often framed on the trends in periodic properties.
  • Conceptual clarity on atomic structure aids in understanding advanced topics in chemistry, physics, and environmental science.
  • Periodic properties help explain real-world phenomena like metal corrosion, formation of compounds, and nuclear reactions.

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