Chapter 8: Structure of Atom

Introduction

• Lecture by Bashir
• Focus on Class 10 syllabus for UP and Telangana state boards.
• Overview of topics covered: quantum numbers, shapes of orbitals, electronic configuration, Pauli's exclusion principle, Aufbau principle, and Hund's rule.

Recap of 9th Standard Learning

• Atoms are spherical (like a football) and mostly hollow.
• Nucleus: Center of the atom, consisting of protons (+ charge) and neutrons (neutral).
• Subatomic Particles: Protons, neutrons, and electrons.
• Electrons revolve around the nucleus like planets around the sun.
• Orbits (or shells):
  • Close to nucleus = less energy.
  • Further from nucleus = more energy.

Orbits and Quantum Numbers

What is an Orbit?

• Defined as the path taken by electrons around the nucleus.
• Orbits are labeled as:
  • First orbit (n=1)
  • Second orbit (n=2)
  • Third orbit (n=3)

Quantum Numbers

• Each electron is described by a set of three quantum numbers: n, l, m.
  • Principal Quantum Number (n): Indicates size and energy of the shell.
    • Higher "n" means larger size and more energy.
    • Examples: n=1 (K shell), n=2 (L shell), n=3 (M shell).
  • Angular Momentum Quantum Number (l): Indicates shape of the subshell; can take values from 0 to n-1.
    • l=0 (s), l=1 (p), l=2 (d), l=3 (f).
  • Magnetic Quantum Number (m): Indicates the orientation of orbitals; values range from -l to +l.

Summary of Quantum Numbers

Maximum Number of Electrons in Subshells

• s subshell: 1 orbital, max 2 electrons.
• p subshell: 3 orbitals, max 6 electrons.
• d subshell: 5 orbitals, max 10 electrons.
• f subshell: 7 orbitals, max 14 electrons.

Electronic Configuration

• Distribution of electrons in shells, subshells, and orbitals.
• Pauli's Exclusion Principle: No two electrons can have the same set of four quantum numbers.
  • Example: Helium atom (2 electrons) both in 1s, but differ in spin.
• Aufbau Principle: Electrons fill orbitals starting with the lowest available energy state (n + l rule).
• Hund's Rule: Electrons will fill degenerate orbitals singly before pairing up.

Examples of Electronic Configuration

• Magnesium (atomic number 12): 1s² 2s² 2p⁶ 3s²
• Chromium (atomic number 24): Anomalies due to stability; configuration: [Ar] 3d⁵ 4s¹.
• Zinc (atomic number 29): Configuration: [Ar] 3d¹⁰ 4s².

Conclusion

• Importance of understanding electronic configuration for predicting chemical behavior.
• Encouragement to interact for questions via website.

Key Points to Remember

1. Atoms are mostly hollow with a dense nucleus made of protons and neutrons.
2. Electrons are arranged in shells around the nucleus, with energy increasing as distance from the nucleus increases.
3. Quantum numbers are crucial for understanding the arrangement of electrons.
4. Use the Aufbau principle, Pauli's exclusion principle, and Hund's rule for determining electronic configuration.
5. Remember exceptions like Chromium and Zinc in electronic configurations for stability.