5.3 Multiple Bonds | Chemistry: Atoms First 2e

Learning Objectives

Describe multiple covalent bonding in terms of atomic orbital overlap
Relate the concept of resonance to π-bonding and electron delocalization

Core Concepts & Principles

While valence bond theory and hybrid orbital models effectively explain molecules with single bonds, they also apply directly to molecules containing double and triple bonds. Multiple covalent bonds are built from combinations of sigma ( $\sigma$ ) and pi ( $\pi$ ) bonds.

Anatomy of Multiple Covalent Bonds

Double Bonds: Consist of one $\sigma$ bond (formed by end-to-end overlap of hybrid orbitals) and one $\pi$ bond (formed by side-by-side overlap of unhybridized parallel $p$ orbitals above and below the internuclear axis).
Triple Bonds: Consist of one $\sigma$ bond and two perpendicular $\pi$ bonds formed by two sets of unhybridized $p$ orbitals.
Bond Rotation: Single ( $\sigma$ ) bonds allow free rotation because their cylindrical symmetry maintains orbital overlap. Multiple bonds restrict rotation because twisting breaks the off-axis side-by-side overlap of the $\pi$ bonds.

Hybridization in Multiple Bonds

$sp^2$ Hybridization (Double Bonds): In molecules like ethene ( $\text{C}_2\text{H}_4$ ), each carbon atom forms three bonding regions (trigonal planar geometry), utilizing three $sp^2$ hybrid orbitals for $\sigma$ bonds. The remaining unhybridized $2p$ orbital on each carbon overlaps side-by-side to form the $\pi$ bond.
$sp$ Hybridization (Triple Bonds): In molecules like acetylene ( $\text{C}_2\text{H}_2$ ), carbon atoms adopt a linear geometry using two $sp$ hybrid orbitals for $\sigma$ bonds. Two pairs of unhybridized $p$ orbitals overlap perpendicularly to form two $\pi$ bonds.

Resonance and Delocalization

Resonance structures involve different arrangements of $\pi$ bonds formed by unhybridized orbitals. Because hybridization depends strictly on steric regions ( $\sigma$ bonds and lone pairs), resonance does not change an atom's hybridization assignment. For example, every carbon atom in benzene ( $\text{C}_6\text{H}_6$ ) is $sp^2$ hybridized regardless of resonance forms. The $\pi$ electrons are not locked in single positions but are instead delocalized throughout the ring system.

Problem-Solving Routines & Methods

Steric Number & Hybridization

\text{Steric Number} = \text{Sigma Bonds} + \text{Lone Pairs}

Determines central atom hybridization by counting effective electron domains.

Variables & Constants

\text{Steric 2}

=sp;

\text{Steric 3}

=sp²;

\text{Steric 4}

=sp³

How to Determine Hybridization in Molecules with Multiple Bonds

1
Draw the Lewis structure of the molecule, noting any resonance forms if present.
2
Count the steric number for the atom of interest: count single bonds, multiple bonds (a double or triple bond counts as ONE effective electron region), and lone pairs.
3
Assign the corresponding hybrid orbital set: 2 regions = sp, 3 regions = sp², 4 regions = sp³.

Pro-Tip: Multiple bonds (double or triple) count as only a single region of electron-pair geometry and hybridization.

Practice & Concept Checks

Concept Check

What is the hybridization and electron-pair geometry of the sulfur atom in sulfur dioxide (

\text{SO}_2

), based on its resonance structures?

Concept Check

Why does rotation occur easily around a single carbon-carbon bond, but is restricted around a carbon-carbon double bond?

Key Terms & Vocabulary

Sigma Bond (σ Bond)Bonding Theory

A covalent bond formed by the end-to-end overlap of atomic orbitals, featuring electron density concentrated symmetrically along the internuclear axis.

Example: C–H single bonds in ethene

Pi Bond (π Bond)Bonding Theory

A covalent bond formed by the side-by-side overlap of unhybridized parallel p orbitals, with electron density located in lobes above and below the internuclear axis.

Example: The second bond in a carbon-carbon double bond

sp² HybridizationHybridization

The mixing of one s and two p atomic orbitals to create three equivalent hybrid orbitals directed in a trigonal planar arrangement.

Example: Carbon atoms in ethene (C₂H₄)

sp HybridizationHybridization

The mixing of one s and one p atomic orbitals to create two equivalent hybrid orbitals directed in a linear arrangement.

Example: Carbon atoms in acetylene (C₂H₂)

Unhybridized p OrbitalAtomic Orbitals

An atomic p orbital that does not participate in hybridization, remaining perpendicular to the hybrid orbital plane to participate in π bonding.

Example: The remaining 2p orbital on carbon in ethene

ResonanceMolecular Structure

A condition where two or more valid Lewis structures with identical atomic frameworks but different π-electron placements describe a single molecule.

Example: Alternate double bond arrangements in benzene

DelocalizationBonding Theory

The sharing of π electrons across three or more adjacent atoms rather than being localized between a specific pair of bonded atoms.

Example: π electrons distributed across the entire benzene ring

Ethene (C₂H₄)Molecules

A simple hydrocarbon featuring a carbon-carbon double bond and sp² hybridized carbon atoms.

Example: C₂H₄

Acetylene (C₂H₂)Molecules

A linear hydrocarbon featuring a carbon-carbon triple bond and sp hybridized carbon atoms.

Example: HC≡CH

Benzene (C₆H₆)Molecules

A cyclic hydrocarbon with a planar hexagonal ring structure where π electrons are fully delocalized through resonance.

Example: C₆H₆