Lesson Notes By Weeks and Term v4 - SHS 2
BONDING
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BONDING
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Subject: Chemistry
Class: SHS 2
Term: 2nd Term
Week: 12
Grade code: 2.2.2.LI.2
Strand code: 2
Sub-strand code: 2
Content standard code: 2.2.2.CS.1
Indicator code: 2.2.2.LI.2
Theme: SYSTEMATIC CHEMISTRY OF THE ELEMENTS
Subtheme: BONDING
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This lesson introduces the concept of hybridization, a crucial idea that helps us understand the true shapes of molecules. We have previously learned about how atoms share electrons to form covalent bonds. However, simple orbital theory doesn't fully explain why a molecule like methane (CH₄) has four identical bonds arranged in a perfect tetrahedron. Hybridization is the "missing piece" of the puzzle. It explains how atomic orbitals mix to form new, special orbitals that allow for stronger bonds and specific molecular shapes.
Part 1: The Problem - Why Do We Need Hybridization?
Let's look at Carbon (C), atomic number 6. Ground state electron configuration: 1s² 2s² 2p² Valence electrons are in the 2s and 2p orbitals. The orbital diagram for the valence shell is:
``` ▲▼ ▲ ▲ ------ ----- ----- ----- 2s 2px 2py 2pz ``` This configuration suggests that Carbon should form only two covalent bonds, using the two half-filled 2p orbitals. However, we know that in methane (CH₄), carbon forms four identical, single covalent bonds with four hydrogen atoms. The shape is tetrahedral with bond angles of 109.5°.
How is this possible? The answer is hybridization. Part 2: What is Hybridization?