Chemical Bonding II: Lewis Structures, Molecular Shapes, and Valence Bond Theory

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Chapter 5: Chemical Bonding I – Lewis Structures and Molecular Shapes

Bond Polarity and Electronegativity

Bond polarity describes the distribution of electrical charge between two bonded atoms. It is determined by the difference in electronegativity, which is the ability of an atom to attract shared electrons in a chemical bond.

Electronegativity Trend: Increases across a period (left to right) and decreases down a group in the periodic table.
Bond Polarity: The greater the difference in electronegativity between two atoms, the more polar the bond.
Example: The H–F bond is highly polar because fluorine is much more electronegative than hydrogen.

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist.

Drawing Lewis Structures: Count total valence electrons, arrange atoms, connect with single bonds, complete octets, and assign remaining electrons as lone pairs.
Resonance Structures: When more than one valid Lewis structure can be drawn for a molecule, these are called resonance structures.
Formal Charge: Used to determine the most stable Lewis structure. Calculated as:
Octet Rule Exceptions: Some atoms may have less (e.g., H, B) or more (expanded octet, e.g., P, S) than 8 valence electrons.
Example: The nitrate ion (NO3-) has three resonance structures.

Bond Strength and Length

The strength and length of a bond depend on the number of shared electron pairs.

Single Bond: One pair of electrons; longest and weakest.
Double Bond: Two pairs of electrons; shorter and stronger than single bonds.
Triple Bond: Three pairs of electrons; shortest and strongest.
Example: In N2, the triple bond makes the molecule very stable and short.

VSEPR Theory: Shapes and Bond Angles

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape of molecules based on electron pair repulsion.

Electron Domains: Regions of electron density (bonds or lone pairs) around a central atom.
Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.
Lone Pairs: Lone pairs occupy more space and can reduce bond angles.
Example: In NH3, the lone pair on nitrogen reduces the H–N–H bond angle from 109.5° to about 107°.

Molecular Shape and Polarity

The shape of a molecule and the distribution of polar bonds determine the overall molecular polarity.

Nonpolar Molecules: Symmetrical shape with equal bond polarities (e.g., CO2).
Polar Molecules: Asymmetrical shape or unequal bond polarities (e.g., H2O).
Example: Water is polar due to its bent shape and polar O–H bonds.

Chapter 6: Chemical Bonding II – Valence Bond and Molecular Orbital Theory

Hybridization of Atomic Orbitals

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds.

Types of Hybridization: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), sp3d, sp3d2.
Example: In methane (CH4), carbon is sp3 hybridized.

Sigma (σ) and Pi (π) Bonds

Bonds between atoms can be classified as sigma (σ) or pi (π) bonds.

Sigma Bond (σ): Formed by head-on overlap of orbitals; all single bonds are sigma bonds.
Pi Bond (π): Formed by side-to-side overlap; present in double and triple bonds (in addition to a sigma bond).
Example: Ethene (C2H4) has a double bond consisting of one sigma and one pi bond.

Molecular Orbital (MO) Theory

Molecular orbital theory describes the distribution of electrons in molecules in terms of molecular orbitals that can extend over the entire molecule.

Bonding and Antibonding Orbitals: Constructive overlap forms bonding orbitals (lower energy), destructive overlap forms antibonding orbitals (higher energy).
MO Diagrams: Used to show the relative energy and occupancy of molecular orbitals in diatomic molecules.
Example: In O2, the presence of unpaired electrons in antibonding orbitals explains its paramagnetism.

Bond Order and Magnetic Behavior

Bond order indicates the strength and stability of a bond, while the presence of unpaired electrons determines magnetic properties.

Bond Order Formula:
Paramagnetic: Molecules with unpaired electrons (attracted to magnetic fields).
Diamagnetic: Molecules with all electrons paired (repelled by magnetic fields).
Example: O2 is paramagnetic; N2 is diamagnetic.

Summary Table: Bond Types and Properties

Bond Type	Number of Shared Electron Pairs	Bond Strength	Bond Length
Single	1	Weakest	Longest
Double	2	Intermediate	Intermediate
Triple	3	Strongest	Shortest