Atomic Structure

Grade 10 · Chemistry

Semester 1 | Period 2 | Week 11

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Subject: Chemistry

Semester: 1

Period: 2

Week: 11

School Name:
Teacher’s Name:
Subject: Chemistry
Grade Level: Grade 10
Week & Period: Week 11, Period II
Date:
Topic: Atomic Structure
Sub-topic: Electron Configuration and Rules/Principles for Electron Filling

 

Learning Objectives

By the end of the lesson, learners should be able to:

  1. Write electron configurations in different formats: full notation, noble gas notation, dot notation, and orbital diagram.
  2. Identify and apply the three major rules for electron filling: Aufbau principle, Pauli Exclusion Principle, and Hund’s Rule.
  3. Use the periodic table to determine the configuration of elements.
  4. Explain the arrangement of electrons across energy levels and sublevels.

 

Previous Knowledge

Learners have studied the structure of the atom, quantum numbers, and types of orbitals (s, p, d, f). They understand that electrons occupy orbitals with specific shapes and energy levels.

 

Instructional Materials

  • Periodic table
  • Orbital filling diagrams
  • Chart of electron configuration formats
  • Flashcards with atomic numbers of common elements
  • Colored pencils for orbital box diagrams

 

Anticipation (Warm-Up) – 5 minutes

Ask:

  • “Why doesn’t an element like oxygen have all its electrons in one shell?”
  • “What do you think determines how many electrons go into each energy level or orbital?”

Introduce today’s lesson as one that explains the rules that control how electrons are arranged in an atom.

 

Building Knowledge (Main Lesson) – 25 minutes

  1. Review of Orbitals
    • s = 2 electrons
    • p = 6 electrons
    • d = 10 electrons
    • f = 14 electrons
  2. Electron Configuration Formats
    Full configuration (e.g., 1s² 2s² 2p⁶...)
    b. Noble gas configuration (e.g., [He] 2s² 2p⁶)
    c. Orbital box diagrams (↑↓ in each box)
    d. Dot notation (used for valence electrons)
  3. Rules for Electron Filling
    Aufbau Principle – Electrons fill from lower to higher energy levels.
    b. Pauli Exclusion Principle – No two electrons in an atom can have the same set of four quantum numbers.
    c. Hund’s Rule – Electrons occupy orbitals singly before pairing.
  4. Examples
    • Hydrogen (1s¹), Carbon (1s² 2s² 2p²), Neon (1s² 2s² 2p⁶), Magnesium ([Ne] 3s²)
    • Use diagrams to represent these configurations

 

Learners’ Activities

  • Write full and noble gas configurations for elements 1–20
  • Draw orbital box diagrams for oxygen, nitrogen, and fluorine
  • Complete worksheets with missing electrons or wrong configurations to correct
  • Group work: create element cards that display atomic number, electron configuration, and orbital diagrams

 

Consolidation (Review and Assessment) – 10 minutes

Oral Questions:

  • What rule tells us to fill orbitals singly before pairing?
  • Which orbital is filled after 3p?
  • What does the configuration [Ne] 3s² represent?

Homework / Assignment:

  1. Draw the orbital box diagrams for atoms of Na, Cl, and Ca
  2. Write full and noble gas electron configurations for K, S, Al
  3. Create a table showing 5 elements, their atomic numbers, and all configurations

 

Notes – Detailed and Explained

  • Electron configuration shows the arrangement of electrons around the nucleus of an atom.
  • There are 4 orbitals: s (2 electrons), p (6), d (10), and f (14).
  • Aufbau Principle: electrons fill orbitals in order of increasing energy (1s → 2s → 2p → 3s...).
  • Pauli Exclusion Principle: no two electrons can have the same four quantum numbers.
  • Hund’s Rule: in a set of orbitals, electrons fill singly before they pair.
  • Configurations can be written in full or shortened using noble gases.
  • Dot notation shows only the valence (outermost) electrons.

 

Expanded Notes / Instructions

  • Reinforce the idea that electron configuration explains chemical behavior of elements.
  • Allow learners to use colored pens to separate energy levels for clarity.
  • Use real-life examples: explain why Na and Cl combine based on their valence electrons.
  • Offer extra guidance for struggling learners when writing noble gas configurations.

 

Inclusive / Differentiation

  • Diagrams and physical models for visual and kinesthetic learners
  • Group problem solving for collaborative learners
  • Step-by-step scaffolding of examples for those who need extra support
  • Encourage peer teaching for students who excel in this area

 

Teacher’s Reflection (Post-Lesson Questions)

  • Did learners understand how the 3 rules guide electron configuration?
  • Were they able to move between full, noble gas, and orbital notations confidently?
  • Which learners struggled with configuration transitions (e.g., from 3p to 4s)?
  • Did my activities reach all types of learners effectively?
  • Should I revise and repeat part of this topic in the next review session?