Hybrid Orbitals and Molecular Shape: VSEPR Theory and Hybridization

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Hybrid Orbitals & Molecular Shape

Introduction

This section explores the concept of hybrid orbitals, their formation, and their role in determining molecular shape. The Valence Shell Electron Pair Repulsion (VSEPR) theory is introduced as a tool for predicting molecular geometry, with a focus on carbon as a model element for hybridization.

Intended Learning Outcomes

Describe atomic orbitals and explain how they combine to generate hybrid orbitals.
Predict the shape and orientation in space of hybrid orbitals.
Arrange hybrid orbitals in order of relative energy.
Apply VSEPR theory to predict the shape of molecules.

VSEPR Theory

Principles of VSEPR

VSEPR (Valence Shell Electron Pair Repulsion) theory states that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion. This determines the geometry of the molecule.

Bonding pairs and non-bonding pairs (lone pairs) both contribute to electron repulsion, but lone pairs exert greater repulsive force, distorting bond angles.
Steric number: The sum of the number of atoms bonded to a central atom and the number of lone pairs on the central atom.

Examples: Water (H2O), Ammonia (NH3), Methane (CH4), Carbon dioxide (CO2).

Steric Number	Lone Pairs	Geometry	Example
2	0	Linear	CO2
3	0	Trigonal planar	BF3
3	1	Bent	SO2
4	0	Tetrahedral	CH4
4	1	Trigonal pyramidal	NH3
4	2	Bent	H2O

The Strange Case of Carbon

Electronic Configuration and Bonding

Carbon has the atomic number 6 and the ground state electron configuration:

In its ground state, carbon has only two unpaired electrons, suggesting it should form only two bonds.
However, carbon commonly forms four bonds (e.g., in methane, CH4), which requires an explanation.

Hybrid Orbitals

Definition and Formation

Hybridization is the process by which atomic orbitals mix to form new, equivalent hybrid orbitals suitable for the pairing of electrons to form chemical bonds in molecules.

For carbon, one 2s electron is promoted to the empty 2p orbital, resulting in four unpaired electrons.
These four orbitals (one 2s and three 2p) combine to form four equivalent sp3 hybrid orbitals.
Each hybrid orbital forms a sigma (σ) bond with another atom, allowing carbon to form four covalent bonds.

Hybrid Orbitals: Relative Energy

Types of Hybridization

Mixing atomic orbitals of different energy can lead to three main types of hybrid orbitals:

sp3 hybridization: Mixing one s and three p orbitals (tetrahedral geometry, 25% s and 75% p character).
sp2 hybridization: Mixing one s and two p orbitals (trigonal planar geometry, 33% s and 67% p character).
sp hybridization: Mixing one s and one p orbital (linear geometry, 50% s and 50% p character).

Hybridization	Orbitals Mixed	Geometry	s Character	p Character
sp3	1 s + 3 p	Tetrahedral	25%	75%
sp2	1 s + 2 p	Trigonal planar	33%	67%
sp	1 s + 1 p	Linear	50%	50%

Hybrid Orbitals: Shape & Orientation

sp3 Hybridization (Methane, CH4)

Four sp3 hybrid orbitals are arranged in a tetrahedral geometry with bond angles of 109.5°.
Each orbital has two lobes of different size; the larger lobe participates in bonding.
Example: Methane (CH4)

sp2 Hybridization (Ethene, C2H4)

Three sp2 hybrid orbitals are arranged in a trigonal planar geometry with bond angles of 120°.
The unhybridized p orbital is perpendicular to the plane and forms a π bond.
Example: Ethene (C2H4)

sp Hybridization (Acetylene, C2H2)

Two sp hybrid orbitals are arranged linearly with a bond angle of 180°.
The two unhybridized p orbitals form two π bonds (triple bond).
Example: Acetylene (C2H2)

Molecular Shape & Hybrid Orbitals

Summary

The arrangement of valence electrons as predicted by VSEPR theory is a direct result of the hybridization of atomic orbitals.
Hybrid orbitals explain the observed shapes and bond angles in organic molecules.