Molecular Geometry and Bonding
Theories
9.1 Molecular Shapes
The size and shape of a molecule of a particular substance play an important part in determining the physical and chemical
properties of that substance. Although Lewis structures are quite useful, they do not indicate a molecule's shape. In this section
the Lewis structure for methane will be compared with other representations of the methane molecule.
9.2 The VSEPR Model
Based on repulsion between like charges, the VSEPR model helps us account for the arrangements of atoms in polyatomic
molecules. The term electron domain is introduced, and you will see an animation that details the arrangements of electron
domains about a central atom.
9.3 Polarity of Polyatomic Molecules
Polarity is revisited in the context of polyatomic molecules. A simulation will show how to determine the net dipole moment, if
any, of a molecule containing polar bonds.
9.4 Covalent Bonding and Orbital Overlap
The bonding in simple molecules such as H2 and Cl2 is explained by overlap of singly occupied atomic orbitals. This is the
introduction to valence-bond theory.
9.5 Hybrid Orbitals
Simple valence-bond theory fails to explain how atoms with no unpaired electrons can form covalent bonds. The use of hybrid
orbitals shows us how bonds can form in such molecules as BeF2 and BCl3. An animation illustrates the hybridization of nitrogen
orbitals in ammonia.
9.6 Multiple Bonds
Lewis structures sometimes contain double or triple bonds. In this section you will see how the bonds are formed by a
combination of hybrid orbitals and unhybridized atomic orbitals.
9.7 Molecular Orbitals
Certain properties of molecules cannot be explained with valence-bond theory or with orbital hybridization. A new theory is
presented to explain some of those properties. You will learn the term bond order and how to calculate it as a measure of a
bond's stability.
9.8 Second-Row Diatomic Molecules
The molecular orbital energy diagrams for homonuclear diatomic molecules are given. The ordering of orbital energies is
explained in terms of orbital interaction. The MO energy-level diagram for oxygen predicts that it should have two unpaired
electrons and should be paramagnetic, which it is. This is one of the things that valence-bond theory could not explain. The
progression from Lewis structures to VSEPR to valence-bond theory to hybridization and finally to molecular orbital theory is an
outstanding example of the scientific method at work.
Molecular Shapes
• Valence-shell electron-pair repulsion (VSEPR)
model
• The molecular geometry of a molecule (or
ion) refers to the arrangement of the atoms.
• The term electron domain refers to a lone pair
or a bond—whether that bond is single,
double, or triple
Polarity of Polyatomic Molecules
The bond dipole is the dipole
moment due only to the two
atoms bonded together. The bond
dipoles are equal in magnitude,
yet exactly opposite one another,
and they cancel each other out.
The overall dipole moment is zero.
Covalent Bonding and Orbital Overlap
• In valence-bond theory we envision atomic
orbitals overlapping as atoms approach one
another, which corresponds nicely to the VSEPR
model. The region of orbital overlap provides
the space in which bonding electron pairs reside.
Hybrid Orbitals
Multiple Bonds
• The single bonds are called sigma (σ) bonds. In a bond the
shared electron density lies directly along the axis between
the two bonded nuclei (called the internuclear axis). A
double bond actually consists of a s bond and a bond. A pi
(π) bond is a covalent bond in which the overlap regions lie
above and below the internuclear axis. Unlike a bond, in
a bond there is no probability of finding the electron on
the internuclear axis; in other words, the electron density
does not lie directly between the atoms.
Single Bond
Pi bond (double bond)
Triple Bond
Molecular Orbitals
• Some aspects of bonding are better explained
by another model called the molecular orbital
theory.
• Atomic orbitals overlap to form molecular
orbitals. When two atomic orbitals of equal
energy overlap, they produce two molecular
orbitals.
Molecular Orbitals (continued)
• A measure of the strength (and stability) of a
bond in a diatomic molecule is the bond
order. Bond order is determined by
Second-Row Diatomic Molecules
• The number of MOs formed equals the number of atomic
orbitals combined.
• Atomic orbitals combine most effectively with other atomic
orbitals of similar energy.
• The effectiveness with which two atomic orbitals combine
is proportional to their overlap with one another; that is, as
the overlap increases, the bonding MO is lowered in energy
and the antibonding MO is raised in energy.
• Each molecular orbital can accommodate, at most, two
electrons, with their spins paired (Pauli exclusion principle).
• When MOs have the same energy, one electron enters each
orbital (with the same spin) before spin pairing occurs
(Hund's rule).
Unpaired electrons cause a molecule to be attracted to a magnetic
field, paramagnetism. Molecules without unpaired electrons are weakly repelled
from a magnetic field, and this type of magnetic behavior is
called diamagnetism.