Term: Second Term
Week: Week 6
Class: SSS 1 (Senior Secondary School 1)
Subject: Chemistry
Topic: Structure of the Atom and Electronic Configuration
Sub-Topics:
Bohr’s Model of the Atom (Revision and Core Assumptions)
Rules for Electronic Configurations (s, p, d, f Orbitals introduction)
1. Bohr’s Model of the Atom (Review)
As we built upon in the first term, the atom is not solid but consists of a central, positively charged nucleus surrounded by moving electrons. The planetary arrangement we use to visualize this comes directly from Niels Bohr.
Core Assumptions (Postulates)
Energy Levels: Electrons move only in specific circular paths called shells or energy levels designated by numbers (n = 1, 2, 3...) or letters (K, L, M,N, O, P, Q)
Quantized States: While an electron stays in its ground-state shell, it does not lose or gain energy.
Quantum Jumps: An electron can jump to a higher, excited shell only by absorbing a fixed packet of energy (a quantum). When it falls back down, it emits that energy as light.
2. Refining Electronic Configuration: Beyond K, L, M...
In the first term, we used the basic 2n^2 rule to distribute electrons into shells (K=2, L=8, M=8 or 18). However, as we look deeper into atomic structure in the second term, we discover that main energy levels (shells) are divided into sub-shells, which are further split into orbitals.
An orbital is a region in space around the nucleus where there is a high probability (greater than 90%) of finding an electron.
The Four Sub-shells (s, p, d, f)
Each primary shell contains a specific number of sub-shells:
| Sub-shell | Number of Orbitals | Max Electron Capacity | Shape |
| s | 1 orbital | 2 electrons | Spherical |
| p | 3 orbitals | 6 electrons | Dumbbell-shaped |
| d | 5 orbitals | 10 electrons | Cloverleaf / Complex |
| f | 7 orbitals | 14 electrons | Highly Complex |
3. Fundamental Rules for Electronic Configuration
To arrange electrons accurately into these sub-shells, chemists follow three strict principles:
A. The Aufbau Principle
Aufbau is a German word meaning "building up."
The principle states that electrons must occupy the orbitals of lowest energy first before moving to higher energy levels.
The energy order does not always follow strict numerical order (for example, the 4s sub-shell is lower in energy than the 3d sub-shell). To remember the order, we use the Aufbau Diagonal Diagram:
B. Pauli’s Exclusion Principle
This principle states that an orbital can hold a maximum of two electrons, and they must have opposite spins. If one electron spins clockwise (↑), the second electron in that same orbital must spin counter-clockwise (↓).
C. Hund’s Rule of Maximum Multiplicity
This rule applies to sub-shells with multiple orbitals of identical energy (like p or d). It states that orbitals are occupied singly by electrons with parallel spins before they begin to pair up.
Think of it like passengers boarding a bus: everyone prefers their own double seat first before pairing up with a stranger.
4. Worked Examples (Orbital Notation)
Let's write out the sub-shell configurations for a few key elements from the first 20:
Oxygen (O, Atomic Number = 8):
Fill 1s first (2 electrons) →1S2
Fill 2s next (2 electrons) → 2S2
Put remaining 4 into 2p → 2P4
Configuration: 1S2 2S2 2P4
Sodium (Na, Atomic Number = 11):
Fill 1s, 2s, and 2p completely 1S2 2S2 2P6(This accounts for 10 electrons).
Put the last electron into 3s → 3s^1
Configuration: 1s^2 2s^2 2p^6 3s^1
Calcium (Ca, Atomic Number = 20):
Following the Aufbau path: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
(Note how 4s fills completely before 3d starts).
Classwork / Evaluation Exercises
What is an atomic orbital? How does it differ from Bohr's idea of an "orbit"?
State Hund’s Rule of Maximum Multiplicity and explain how it applies when placing 3 electrons into the 2p sub-shell.
Write the s, p, d, f orbital electronic configurations for:
Nitrogen (Atomic Number = 7)
Magnesium (Atomic Number = 12)
Chlorine (Atomic Number = 17)
Why does the 4s sub-shell fill up before electrons enter the 3d sub-shell?
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