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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:

Chapter 9 - Molecular Geometry and Bonding Theories

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) electron-

domain 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 valence-

bond 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