This lesson explains how electrons are arranged in atoms, how orbitals fill, and how these arrangements determine trends in the periodic table.

1. Core Concepts (Short Notes)

4.1 Electronic Configuration

Electronic configuration describes the arrangement of electrons in shells, subshells, and orbitals.

• Shells: Represented by n = 1, 2, 3…

• Subshells: s, p, d, f

• Orbitals: Regions where electrons are likely to be found.

Each orbital holds 2 electrons.

4.2 Aufbau Principle

Electrons fill orbitals in order of increasing energy. Order (simplified): 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → …

4.3 Hund’s Rule

Electrons fill each orbital singly before pairing.

• Helps minimize electron–electron repulsion.

  
  

4.4 Pauli’s Exclusion Principle

No two electrons in the same atom can have the same 4 quantum numbers.

• Means: each orbital can hold only 2 electrons with opposite spins.

  
  

4.5 Short and Full Notation

• Full: 1s² 2s² 2p⁶

• Shorthand: [Ne] 3s² 3p⁴

  
  

4.6 Periodicity (Periodic Trends)

The properties of elements change across a period or down a group due to electron arrangement.

Atomic Radius

• Decreases across a period (more protons → stronger pull).

• Increases down a group (more shells added).

Ionization Energy

Energy required to remove an electron.

• Increases across a period.

• Decreases down a group.

Electronegativity

Ability of an atom to attract electrons.

• Increases across a period.

• Decreases down a group.

Electron Affinity

Energy change when an atom gains an electron.

• Similar trend to electronegativity.

  
  

4.7 Why These Trends Occur

Two key factors:

• Nuclear charge: More protons → stronger attraction.

• Shielding effect: Inner electrons repel outer electrons.

2. Key Formulas to Memorize

There are no strict formulas in this lesson, but remember:

• Maximum electrons in a shell: 2n²

• Orbitals per subshell: s = 1, p = 3, d = 5, f = 7

3. Tips & Tricks for Exams

• 4s fills before 3d, but 4s electrons are removed first during ion formation.

• For shorthand configuration, start from the nearest noble gas.

• Use the diagonal rule (Aufbau diagram) to fill orbitals easily.

• Look out for exceptions: Cr and Cu show 3d⁵4s¹ and 3d¹⁰4s¹ due to stability.

• For periodic trends, always think: more protons vs. more shells.

• Noble gases have full outer shells → stable and unreactive.

4. Important Points to Remember

• Electronic configuration determines chemical behavior.

• Periodic trends arise from predictable changes in structure.

• Ionization energy drops slightly at group 2→3 and group 5→6.

• Atomic radius changes affect reactivity (e.g., alkali metals vs. halogens).

Elements in the same group have similar chemical properties.