Electronic Configuration – KVPY Chemistry Study Notes
KVPY (Kishore Vaigyanik Protsahan Yojana) – Historic syllabus – Chemistry – Topic: Electronic Configuration
Goal: Understand how electrons are arranged in atoms, why the arrangement matters, and how to write and interpret electronic configurations.
1. Simple Explanation
An atom’s electrons occupy “shells” (energy levels) around the nucleus.
- Shells are labeled with principal quantum number n (1, 2, 3 …).
- Within each shell, sub‑shells (orbitals) are labeled s, p, d, f (corresponding to angular momentum quantum number l = 0, 1, 2, 3).
- Electrons fill these orbitals from lowest to highest energy – this is the electronic configuration of the atom.
Key Idea: The arrangement of electrons determines an element’s chemical properties and its place in the periodic table.
2. Key Points
| Concept | What it Means | Why It Matters |
|---|---|---|
| Principal Quantum Number (n) | Integer 1, 2, 3…; size and energy of shell | Determines the period in the periodic table |
| Azimuthal Quantum Number (l) | 0–(n–1); type of sub‑shell (s, p, d, f) | Determines shape of orbital and electron capacity |
| Magnetic Quantum Number (m<sub>l</sub>) | –l … +l | Number of orbitals in a sub‑shell |
| Spin Quantum Number (m<sub>s</sub>) | ±½ | Pauli Exclusion Principle (two electrons per orbital, opposite spins) |
| Aufbau Principle | Electrons occupy lowest energy orbitals first | Predicts ground‑state configuration |
| Pauli Exclusion Principle | No two electrons in an atom can have identical set of four quantum numbers | Limits electron count per orbital |
| Hund’s Rule | Electrons occupy separate orbitals of the same energy before pairing | Maximizes total spin → lower energy |
| Sub‑shell Capacity | s = 2, p = 6, d = 10, f = 14 | Determines how many electrons can fill each sub‑shell |
| Total Electron Count | Sum of all electrons across sub‑shells | Equals atomic number Z (for neutral atoms) |
| Notation | 1s² 2s² 2p⁶ … or shorthand [Ar]3d⁵ 4s² | Compact way to write configuration |
| Sub‑shell Ordering | 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s … (except for transition metals, lanthanides, actinides) | Shows energy trend (diatomic rule) |
3. Worked Example
Find the electronic configuration of the element with atomic number 29 (Copper, Cu).
Step‑by‑Step
| Step | Action | Result |
|---|---|---|
| 1 | Start from 1s: 1s² (2 e⁻) | 2 |
| 2 | Add 2s: 2s² (2 e⁻) | 4 |
| 3 | Add 2p: 2p⁶ (6 e⁻) | 10 |
| 4 | Add 3s: 3s² (2 e⁻) | 12 |
| 5 | Add 3p: 3p⁶ (6 e⁻) | 18 |
| 6 | Add 4s: 4s² (2 e⁻) | 20 |
| 7 | Add 3d: 3d¹⁰ (10 e⁻) | 30 |
We have 30 electrons after filling 3d, but Cu has only 29.
So the last electron does not go into 3d but into 4s.
However, the actual ground‑state configuration is:
Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
Why? 4s² has a slightly higher energy than 3d, but when 3d is filled, 4s becomes lower. The electron prefers 4s¹, leaving 3d with 10 electrons.
Compact (Short) Notation
Cu: [Ar] 3d¹⁰ 4s¹
where [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶.
4. Quick Revision Summary
- Electron arrangement → governs chemical behavior.
- Quantum numbers: n, l, m<sub>l</sub>, m<sub>s</sub> define each electron’s state.
- Rules: Aufbau → Pauli → Hund; sub‑shell capacity = 2(2l+1).
- Order of filling: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s …
- Shorthand: Use noble‑gas core + valence configuration.
- Transition metals, lanthanides, actinides: 4s/3d, 5s/4f, 6s/5f interactions create exceptions.
- Check: Sum of electrons = atomic number.
- Practice: Write configurations for elements in each block (s, p, d, f) and identify noble‑gas cores.
Remember: Mastering electronic configuration is the gateway to understanding periodic trends, bonding, magnetism, and spectral properties. Happy studying!