chemical bonding chemistry lesson notes
chemical bonding chemistry lesson notes
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Topic 4: Chemical Bonding, Structure and Intermolecular Forces
Topic Objectives
By the end of this topic, students should be able to:
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Explain the electronic theory of bonding.
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Briefly describe the nature and characteristics of ionic, covalent, coordinate, metallic, hydrogen, Van der Waals, and dipole-dipole bonds/forces, giving examples of elements or compounds where they exist.
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Show how electronegativity influences the formation of ionic and covalent bonds.
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Predict and explain the shapes of simple molecules and ions using the Valence Shell Electron Pair Repulsion (VSEPR) theory.
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Indicate and give examples of how bond type influences the physical and chemical properties of elements and compounds.
Lesson Topic: 4.1 – Electronic Theory of Bonding
Lesson Objectives
By the end of this lesson, students should be able to:
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State and explain the electronic theory of bonding.
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Define a chemical bond and name its different types.
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Describe the nature and formation of ionic bonding.
I. Introduction
The Electronic Theory of Bonding
The Electronic Theory of Bonding states that during the formation of a chemical bond, atoms attempt to achieve a stable outer-shell electron configuration similar to that of the nearest noble gas in the Periodic Table. This stable state is reached by either losing, gaining, or sharing valence electrons.
Examples of Electron Loss (Cations):
Metals typically lose their valence electrons to achieve the stable configuration of the next lower noble gas:
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Sodium ($\text{Na}$), Magnesium ($\text{Mg}$), and Aluminium ($\text{Al}$) lose one, two, and three valence electrons respectively.
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They form $\text{Na}^+$, $\text{Mg}^{2+}$, and $\text{Al}^{3+}$ ions, all achieving the stable electronic configuration of Neon ($\text{Ne}$):
Examples of Electron Gain (Anions):
Non-metals typically gain electrons to achieve the stable configuration of the next higher noble gas:
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Sulfur ($\text{S}$) and Chlorine ($\text{Cl}$) gain two and one electrons respectively.
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They form $\text{S}^{2-}$ and $\text{Cl}^-$ ions, achieving the stable electronic configuration of Argon ($\text{Ar}$):
⚠️ Important Exceptions to the Rule
While the octet rule applies broadly, there are significant exceptions where atoms do not attain a strict noble gas structure:
Transition Metal Ions: Most transition metal ions, such as $\text{Fe}^{2+}$, $\text{Fe}^{3+}$, and $\text{Cu}^{2+}$, do not achieve a noble gas configuration upon losing electrons.
Expanded Octets: The sulfur ($\text{S}$) atom in compounds like sulfur dioxide ($\text{SO}_2$) and sulfur trioxide ($\text{SO}_3$) expands its valence shell beyond the standard octet structure.
