Chemistry Lesson Note On Ionic and Covalent Bonding

Topic: Ionic and Covalent Bonding

Class: SS1

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Specific Objectives:

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

1. Define chemical bonding.
2. Distinguish between ionic and covalent bonds.
3. Explain the formation of ionic and covalent bonds with examples.
4. List the physical and chemical properties of ionic and covalent compounds.
5. Draw electron dot (Lewis) structures of simple covalent and ionic compounds.

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Instructional Materials:

• Models showing ionic and covalent bonds
• Periodic table
• Flashcards of metals and non-metals
• Ball and stick models
• Charts showing Lewis structures

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Lesson Content:

Step 1: Introduction to Chemical Bonding

• A chemical bond is the force of attraction that holds atoms together in a compound.
• Atoms bond to achieve a stable octet (8 electrons in their outer shell), similar to noble gases.
• There are two main types of bonding:
  1. Ionic bonding
  2. Covalent bonding

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Step 2: Ionic Bonding (Electrovalent Bond)

• Definition: Ionic bond is formed by the transfer of electrons from one atom (usually a metal) to another (usually a non-metal).
• The metal loses electrons to become a positive ion (cation), and the non-metal gains electrons to become a negative ion (anion).
• The oppositely charged ions are held together by electrostatic attraction.

Examples:

• Sodium chloride (NaCl):

  • Na (2,8,1) loses 1 electron → Na⁺
  • Cl (2,8,7) gains 1 electron → Cl⁻
  • Na⁺ + Cl⁻ → NaCl

• Magnesium oxide (MgO):

  • Mg (2,8,2) → Mg²⁺
  • O (2,6) → O²⁻
  • Mg²⁺ + O²⁻ → MgO

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Step 3: Covalent Bonding (Molecular Bond)

• Definition: Covalent bond is formed when two non-metallic atoms share one or more pairs of electrons to achieve a stable octet.
• The shared electrons belong to both atoms.

Types of Covalent Bonds:

• Single covalent bond: Sharing of 1 pair of electrons (e.g. H₂)
• Double covalent bond: Sharing of 2 pairs of electrons (e.g. O₂)
• Triple covalent bond: Sharing of 3 pairs of electrons (e.g. N₂)

Examples:

• Hydrogen molecule (H₂):
  H • + • H → H:H or H–H

• Water (H₂O):
  Each hydrogen shares 1 electron with oxygen.

• Methane (CH₄):
  Carbon shares one electron each with 4 hydrogen atoms.

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Step 4: Comparison of Ionic and Covalent Compounds

Property	Ionic Compounds	Covalent Compounds
Formation	Transfer of electrons	Sharing of electrons
Types of elements involved	Metals and non-metals	Non-metals only
Structure	Crystal lattice	Molecules
Melting/Boiling point	High	Low (except giant covalent)
Electrical conductivity	Conducts when molten/dissolved	Poor (except in aqueous solutions)
Solubility	Soluble in water	Soluble in non-polar solvents

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Step 5: Drawing Electron Dot Structures (Lewis Structures)

• Ionic compound (NaCl):

  • Na: 1 valence electron → [Na]⁺
  • Cl: 7 valence electrons → [Cl⁻] with 8 electrons

• Covalent compound (H₂O):

  • O shares one electron each with two H atoms → H:O:H

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Evaluation:

1. Define chemical bonding.
2. Differentiate between ionic and covalent bonding with two examples each.
3. State three properties of covalent and ionic compounds.
4. Using electron dot diagrams, show the formation of NaCl and H₂O.
5. Which type of bond is present in: (a) CH₄ (b) MgCl₂ (c) CO₂?

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Homework:

1. Draw the electron dot structure for:
   • Ammonia (NH₃)
   • Carbon dioxide (CO₂)
   • Magnesium chloride (MgCl₂)
2. Write five differences between ionic and covalent bonds in a tabular form.
3. Explain why noble gases do not form chemical bonds.
4. Give two real-life applications of ionic and covalent compounds.

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Conclusion:

Chemical bonds hold atoms together to form compounds. Ionic bonds involve the transfer of electrons, while covalent bonds involve sharing. Understanding these helps explain the behavior and properties of substances in chemistry.