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Chapter 9 - Molecular Geometry and Bonding Theories
9.1 Molecular Shapes
Lewis structures only provide the number and types of bonds
between atoms
Shapes of ABn molecules depend on the value of n. The shapes for
AB2 and AB3 molecules are:
In order to predict molecular shape, we assume that the valence
electrons repel each other
For molecules of the general form ABn there are 5 fundamental
shapes:
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The shape of any particular ABn molecule can be obtained by
removing corner atoms from one of these basic shapes. Beginning
with a tetrahedron:
9.2 The VSEPR Model
What determines the shape of a molecule? Electron pairs, whether
bonding or nonbonding, repel each other.
Each nonbonding pair, single bond, or multiple bond produces an
electron domain around the central atom.
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The shapes of different ABn molecules or ions depend on the number
of electron domains surrounding the central A atom.
We can use the electron-domain geometry to help us predict the
molecular geometry.
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Effect of nonbonding electrons and multiple
bonds on bond angles
A nonbonding pair of electrons are physically
larger than bonding electron pairs.
Electrons in nonbonding pairs and in multiple bonds repel more
than electrons in single bonds:
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Molecules with Expanded Valence Shells
Atoms that have expanded octets have five electron domains
(trigonal bipyramidal) or six electron domains (octahedral) electrondomain geometries.
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Shapes of larger Molecules
We can use the VSEPR model to predict the geometry around interior
atoms in complex molecules.
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Example
Give the electron-domain and molecular geometries for the following
molecules and ions: (a) HCN, (b) SO32–, (c) SeF4, (d) PF6–, (e) BF4–,
(f) N3–.
9.3 Molecular Shapes and Molecular Polarity
Bond polarity  How equally the electrons are shared between the
two atoms
Dipole moment  Amount of charge separation
For a molecule that consists of more than two atoms, the dipole
moment depends on both the polarities of the individual bonds and
the geometry of the molecule.
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9.4 Covalent Bonding and Orbital Overlap
Valence bond theory  combination of Lewis theory and valencebond theory
Lewis theory  covalent bonds occur when atoms share electrons
Valence-bond theory  merging of two valence atomic orbitals from
different atoms leading to a buildup of electron density.
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The potential energy of the system changes as two H atoms come
together to form a H2 bond.
9.5 Hybrid Orbitals
To extend the ideas of orbital overlap and valence-bond theory to
polyatomic molecules, we need to introduce the concept of hybrid
orbitals.
We must explain both the formation of electron-pair bonds and the
observed geometries of the molecules.
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sp Hybrid Orbitals
To illustrate the concept of hybridization we will consider the BeF2
molecule.
Consider beryllium: in its ground electronic state, it is incapable of
forming bonds with fluorine atoms.
The 2s orbital and one of the 2p orbitals can “mix” to generate two new
degenerate orbitals.
These two degenerate orbitals align themselves 180° from each other,
and 90° from the two remaining unhybridized p orbitals.
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sp2 and sp3 Hybrid Orbitals
Three equivalent sp2 hybrid orbitals are formed from hybridization of
one s and two p orbitals.
sp2 and sp3 Hybrid Orbitals
Four equivalent sp3 hybrid orbitals are formed from the hybridization of
one s and all three p orbitals.
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Hybridization Involving d Orbitals
Atoms in the third period and beyond can form hybrid orbitals using d
orbitals
Five equivalent sp3d hybrid orbitals are formed from the hybridization
of one s and three p orbitals and one d orbital.
Six equivalent sp3d2 hybrid orbitals are formed from the hybridization of
one s and three p orbitals and one d orbital.
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Summary
To predict the hydrid orbitals used by an atoms in bonding:
1. Draw Lewis structure.
2. Determine the electron-domain geometry using the VSEPR model.
3. Specify the hydrid orbitals needed to accommodate the electron
pairs based on their geometric arrangement.
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9.6 Multiple Bonds
The covalent bonds we have seen so far are sigma (σ) bonds,
characterized by:
To describe multiple bonding, we must consider a second type of
bond, pi () bonds.
Pi () bonds are characterized by:
Single bonds are always σ bonds, because σ overlap is greater,
resulting in a stronger bond and more energy lowering.
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Example of a double bond - ethylene (C2H4):
In a molecule like formaldehyde an sp2 orbital on carbon overlaps in σ
fashion with the corresponding orbital on the oxygen.
In triple bonds, e.g. acetylene, two sp orbitals form a σ bond between
the carbons
Also possible to make  bonds from d orbitals.
 bonds formed only if unhydridized p orbitals present.
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Delocalized  Bonding
When writing Lewis structures for species like the nitrate ion, we draw
resonance structures to more accurately reflect the structure of the
molecule or ion
The p orbitals on all three oxygens overlap
with the p orbital on the central nitrogen
Another example of delocalized  bonding is benzene (C6H6).
Six electrons are available for  bonding.
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9.7 Molecular Orbitals
Some aspects of bonding are not explained by Lewis structures,
VSEPR theory, and hybridization.
We can use molecular orbital (MO) theory to explain some of these
observations.
Molecular orbitals have some characteristics are similar to those of
atomic orbitals.
The Hydrogen Molecule
When n AOs overlap, n MOs form. For H2, 1s (H) + 1s (H) must result
in 2 MOs:
Both bonding and antibonding molecular orbitals have electron
density centered around the internuclear axis
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Note similarity to mixing of atomic orbitals to make hybrid atomic
orbitals: the difference here is that the orbitals used are on two
different nuclei.
Bond Order
In MO theory the staability of a covalent bond is related to its bond
order.
Bond Order =
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23.5 Metallic Bonding
MO theory helps us explain the bonding in metals. Recall that n AOs
are used to make n MOs.
The energy differences between orbitals are small, so promotion of
electrons requires little energy: electrons readily move through the
metal.
12.1 Metals and Smiconductors
Materials may be classified according to their band structure.
Metals
Good electrical conductors.
Semiconductors
Band structure has an energy gap separating totally filled bands and
empty bands.
Insulators
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