Bonding Theories, Hybridization, and Molecular Orbitals in General Chemistry

Study Guide - Smart Notes

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Bonding Theories in Context

Overview of Bonding Theories

Several bonding theories are used in chemistry to predict and explain the structures and properties of molecules. Each theory has its strengths and limitations, and understanding their differences is crucial for interpreting molecular behavior and reactivity.

Theory	What it does	Failures	Use for
Lewis model	Valence electrons make bonds	Does not explain shape, bond lengths, magnetic properties, resonance/delocalization	Basic structure and connectivity
VSEPR	Uses electron pair repulsion to predict shape and bond angles of molecules	Not a standalone model – use with Lewis model or VBT	Shape and bond angles
VBT (Valence Bond Theory)	Uses atomic or hybrid atomic orbitals to make bonds. Uses VSEPR for shape and angles.	Resonance (delocalization) is not explained well	Hybridization (explains which orbitals contain lone pairs or make bonds)
MO theory (Molecular Orbital Theory)	Combines atomic orbitals to form molecular orbitals. Explains resonance and delocalization.	Too complicated for routine structure and reactivity	Use when all else fails to explain structure and reactivity

Valence Bond Theory (VBT)

Basic Principles

Valence Bond Theory describes covalent bonds as the result of overlapping atomic orbitals. The greater the overlap, the stronger the bond. Atomic orbitals that overlap are typically s, p, d, or hybrid orbitals.

Bond formation: Overlap of half-filled atomic orbitals from two atoms.
Hybridization: Mixing of atomic orbitals to form new, equivalent hybrid orbitals for bonding.
Example: In methane (CH4), carbon forms four equivalent bonds using sp3 hybrid orbitals.

Hybridization

Concept and Need for Hybridization

Hybridization is the process of mixing atomic orbitals (AOs) to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in molecules. This concept explains the observed molecular geometries that cannot be described by simple atomic orbitals alone.

Number of hybrid orbitals formed = Number of atomic orbitals mixed
Hybrid orbitals are degenerate (have the same energy)
Type of hybridization depends on the steric number (number of electron domains around the atom)

Types of Hybridization and Electron Geometry

SN	Hybridization	AOs used	Hybrid AOs formed	Electron geometry	Example
2	sp	1 s + 1 p	2	Linear	BeCl2
3	sp2	1 s + 2 p	3	Trigonal planar	BF3
4	sp3	1 s + 3 p	4	Tetrahedral	CH4
5	sp3d	1 s + 3 p + 1 d	5	Trigonal bipyramidal	PF5
6	sp3d2	1 s + 3 p + 2 d	6	Octahedral	SF6

Practice: Identifying Hybridization

Determine the steric number (SN) for each atom (number of atoms bonded + number of lone pairs)
Assign the corresponding hybridization based on SN

Sigma (σ) and Pi (π) Bonds

Formation of Multiple Bonds

Multiple bonds between atoms consist of one sigma (σ) bond and one or more pi (π) bonds.

σ bond: Head-on overlap of atomic or hybrid orbitals along the internuclear axis. Stronger than π bonds. Allows free rotation around single bonds.
π bond: Sideways overlap of unhybridized p orbitals, above and below the internuclear axis. Restricts rotation around double and triple bonds.

Bond Type	How Formed	Properties
σ bond	Head-on overlap of orbitals	Stronger, allows free rotation
π bond	Sideways overlap of p orbitals	Weaker, restricts rotation

Orbital Overlap Diagrams

Examples: Formaldehyde and Acetylene

Hybrid orbitals are used to make σ bonds or hold lone pairs.
π bonds are made from overlap of unhybridized p orbitals.
In acetylene (C2H2), both C atoms are sp hybridized; H never hybridizes.

Cis-Trans Isomerism

Definition and Consequences

Cis and trans isomers are molecules with the same formula but different spatial arrangements due to restricted rotation around double bonds.

Cis: Substituents on the same side of the double bond
Trans: Substituents on opposite sides
Restricted rotation is due to the presence of π bonds

Molecular Orbital (MO) Theory

Basic Principles

MO theory applies quantum mechanics to molecules, combining atomic orbitals to form molecular orbitals that extend over the entire molecule. This theory explains resonance, delocalization, and magnetic properties.

Bonding MO: Lower energy, more stable, electron density between nuclei
Antibonding MO: Higher energy, less stable, electron density outside the internuclear region
Designated as σ, σ*, π, π*

Bond Order Calculation

Bond order indicates the strength and stability of a bond:

Bond order formula:

Higher bond order = stronger, shorter bond
Bond order of 0 = unstable, molecule does not exist

Magnetic Properties from MO Diagrams

If MO diagram has unpaired electrons: paramagnetic
If all electrons are paired: diamagnetic
MO diagrams can be used to compare stability and magnetic properties of molecules like O2, N2, etc.

Summary Table: Bonding Theories Comparison

Theory	Explains	Fails to Explain	Best Use
Lewis	Connectivity	Shape, resonance, magnetism	Simple molecules
VSEPR	Shape, bond angles	Resonance, delocalization	Predicting geometry
VBT	Hybridization, bond formation	Delocalization	Describing bonds and lone pairs
MO	Resonance, magnetism	Complexity	Advanced explanations

Key Terms and Definitions

Hybridization: Mixing of atomic orbitals to form new hybrid orbitals for bonding
Sigma (σ) bond: Covalent bond formed by head-on overlap of orbitals
Pi (π) bond: Covalent bond formed by sideways overlap of p orbitals
Bond order: Measure of bond strength, calculated from MO theory
Paramagnetic: Substance with unpaired electrons, attracted to magnetic fields
Diamagnetic: Substance with all electrons paired, weakly repelled by magnetic fields

Practice and Application

Identify hybridization for each atom in a molecule by counting electron domains
Determine the number of σ and π bonds in a molecule
Draw MO diagrams and calculate bond order and magnetic properties
Recognize cis-trans isomerism and its relation to π bonds