Chapter 9

Lewis Theory-VSEPR Valence Bond Theory

Molecular Orbital Theory

Problems with Lewis Theory

Lewis theory generally predicts trends in properties, but does not give good numerical predictions.

Lewis theory gives good first approximations of the

bond angles in molecules, but usually cannot be used to get actual bond angles.

Lewis theory cannot write one correct structure for many molecules where resonance is important.

Lewis theory often does not predict the correct magnetic behavior of molecules.

Valence Bond Theory

Linus Pauling and others applied the principles of quantum mechanics to molecules.

They reasoned that bonds between atoms would occur when the atomic orbitals interacted to

make new bonds.

The types of interactions depend on whether the orbitals align along the axis between the nuclei, or

outside the axis.

Orbital Interaction

As two atoms approached, the half-filled valence atomic orbitals on each atom would interact to

form molecular orbitals.

The molecular orbitals would be more stable than the separate atomic orbitals because they would

contain paired electrons shared by both atoms.

Orbital Diagram for the Formation of H2S

S ↑ ↑ ↑↓ ↑↓

3s 3p

↑

H

1s

↑

H

1s

↑↓ H─S bond

↑↓ H─S bond

H S H

Predicts bond angle = 90° Actual bond angle = 92°

“Unhybridized” C Orbitals Predict the Wrong Bonding & Geometry

H 1s

H 1s

C 2s 2p

Valence Bond Theory - Main Concepts

Valence electrons of the atoms in a molecule reside in quantum-mechanical atomic orbitals. The orbitals

can be the standard s, p, d, and f orbitals, or they may be hybrid combinations of these.

A chemical bond results when two of these atomic

orbitals interact and there is a total of two electrons in a new molecular orbital.

The shape of the molecule is determined by the geometry of the interacting orbitals.

Hybridization

Hybridization is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals.

# of new orbitals-----> 2, 3, 4, 5, 6 orbital designation---> sp, sp2, sp3, sp3d, sp3d2

The same type of atom can have different types of hybridization:

C,N = sp, sp2, sp3

The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule.

The sp Hybrid Orbitals in Gaseous BeCl2

Cl Be Cl

The sp2 Hybrid Orbitals in BF3

B

F

FF

The sp3 Hybrid Orbitals in CH4

The sp3 Hybrid Orbitals in NH3

The sp3 Hybrid Orbitals in H2O

Hybridization and VSEPR Theory

All three central atoms are sp3 hybridized !!

The sp3d Hybrid Orbitals in PCl5

The sp3d2 Hybrid Orbitals in SF6

sp3 hybridized ↑ ↑ ↑ ↑ Unhybridized

2s 2p

↑↓ ↑ ↑

Carbon Hybridizations

Unhybridized

2s 2p

↑↓ ↑ ↑ sp2 hybridized

2p 2sp2

↑ ↑ ↑ ↑

Unhybridized

2s 2p

↑↓ ↑ ↑ sp hybridized

2sp

↑ ↑ ↑ ↑

2p

2 sp3

Different Carbon Hybridizations Lead to Different Molecular Geometries

electron density

sp3

sp2

sp

sp3 Hybridization

Atom with four electron groups around it

tetrahedral electron group geometry ~109.5° angles between hybrid orbitals

tetrahedral molecular geometry for carbon trigonal pyramidal geometry for nitrogen

bent geometry for oxygen

Atom uses hybrid orbitals for all bonds & lone pairs

sp3 Hybridized Atoms

Place electrons into hybrid and unhybridized valence orbitals as if all the orbitals have equal energy

CUnhybridized atom

2s 2p

↑↓ ↑

sp3 hybridized atom

2sp3

↑ ↑ ↑ ↑ ↑

N2s 2p

↑↓ ↑

2sp3

↑ ↑ ↑ ↑ ↑ ↑

↑

O↑

↑

2s 2p

↑↓ ↑ ↑ ↑↓

2sp3

↑ ↑ ↑↓

Bonding with Valence Bond Theory

Bonding takes place between atoms when their atomic or hybrid orbitals interact (“overlap”).

To interact, the orbitals must either be aligned along the axis between the atoms, or

The orbitals must be parallel to each other and perpendicular to the interatomic axis.

Bonding in Methane

Ammonia Formation with sp3 N

O

H

H

1s1ssp3

sp3

Water Formation with sp3 O

Types of Bonds

Sigma (σ) bond - when the interacting atomic orbitals point along the axis connecting the two bonding nuclei

Pi (π) bond - when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the

two bonding nuclei. (Usually from overlap of two unhybridized p orbitals)

The interaction between parallel orbitals is not as strong as

between orbitals that point at each other;

Therefore, σ bonds are stronger than π bonds.

Types of Bonds

sp2 Hybridization

Atom with three electron groups around it

trigonal electron group planar system ~120° bond angles - flat

C = trigonal planar molecular geometry N = bent molecular geometry

O = linear geometry

Atom uses hybrid orbitals for σ bonds and lone pairs

Atom uses a nonhybridized p orbital for a π bond

sp2 Hybridized Atoms Orbital Diagrams

Unhybridized atom

2s 2p

↑↓ ↑

sp2 hybridized atom

2sp2

↑ ↑ ↑ ↑ ↑

2p

2s 2p

↑↓ ↑

2sp2

↑ ↑ ↑ ↑

2p

↑ ↑

↑

2s 2p

↑↓ ↑ ↑ ↑↓

2sp2

↑ ↑

2p

↑

↑

↑↓

C

N

O

Formaldehyde, CH2O

↑ sp2 C ↑ ↑

↑ ↑

sp2 O ↑ ↑↓

↑ ↑

1s H

σ

σ σ

1s H

p C p O

π

↑↓

C OH

H

C OH

H

Bonding in Formaldehyde

Hybrid orbitals overlap to form σ bonds. Unhybridized p orbitals overlap to form a π bond.

Ethene, CH2CH2C C

H

H H

H

↑ sp2 C ↑ ↑

↑

↑ ↑

1s H

σ σ

1s H

p C

σ

π

↑ ↑ ↑

↑

↑ ↑

1s H

σ σ

1s H

sp2 C

p C

Bonding in Ethene, C2H4

π

π

Bonding in Ethene

CH2NH Orbital Diagram

↑ sp2 C ↑ ↑

↑ ↑

sp2 N ↑ ↑ ↑↓ ↑ ↑ ↑

1s H

σ

σ σ σ

1s H 1s H

p C p N

C N

H

H H

･ ･

π

C NH

H H

Bond RotationRotation around a σ bond does not require breaking the

interaction between atomic orbitals.

Rotation around a π bond requires the breaking of the interaction between atomic orbitals.

Restricted Rotation Around π-bonded Atoms in C2H2Cl2

Restricted Rotation Around π-bonded Atoms in C2H2Cl2

“cis” “trans”

Restricted Rotation Around π-bonded Atoms in C2H2Cl2

nonet

dipole

sp hybridization

Atom with two electron groups

linear shape 180° bond angle

Atoms use hybrid orbitals for σ bonds or lone pairs

Atom use nonhybridized p orbitals for π bonds

sp Hybridized Atoms Orbital Diagrams

Unhybridized atom

2s 2p

↑↓ ↑

sp hybridized atom

2sp

↑ ↑ ↑ ↑ ↑

2p

2s 2p

↑↓ ↑

2sp

↑ ↑ ↑ ↑

2p

↑ ↑

↑

C

N

C

HCN Orbital Diagram

↑ sp C ↑

↑ ↑ ↑

sp N

↑

↑ ↑↓ ↑

1s H

σ

s

p C p N

2π

C NH

σ

Bonding in HCNH C N

⇵H

⇵

⇵

C

N

HCCH (C2H2) Orbitals C CH H

↑ sp C ↑

↑ ↑

↑

1s H

s

p C σ

2π

↑ ↑

↑ ↑

↑

1s H

s

p C

sp C

σ σ

Bonding in C2H2

Bonding in HCCHH C C H

H

C

C

H

Bonding in C2H2

sp3d hybridization

sp3d hybridization

sp3d hybridization

P

S

Unhybridized atom

3s 3p

↑↓ ↑

sp3d hybridized atom

3sp3d

↑ ↑ ↑ ↑ ↑ ↑

3d

↑

3s 3p

↑↓

3sp3d

↑ ↑ ↑ ↑ ↑ ↑↓

3d

↑↓ ↑

(non-hybridizing d orbitals not shown)

4s 4p

↑↓ ↑↓ ↑↓ ↑

4d 4sp3d

↑↓ ↑ ↑ ↑ ↑↓ Br

sp3d hybridization

sp3d hybridization

sp3d hybridization

sp3d2 hybridization

sp3d2 Hybridized AtomsOrbital Diagrams

S

Br

3s 3p 3sp3d2

↑↓

3d

↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑

Unhybridized atom sp3d2 hybridized atom

4s 4p

↑↓

4d

↑↓ ↑↓ ↑

4sp3d2

↑↓ ↑ ↑ ↑ ↑ ↑

sp3d2 hybridization

sp3d2 hybridization

Atom with six electron groups around it octahedral electron geometry

Square Pyramid, Square Planar 90° bond angles

Use empty d orbitals from valence shell. d orbitals can be used to make π bonds.

Predicting Hybridization and Bonding Scheme

1. Start by drawing the Lewis structure 2. Use VSEPR Theory to predict the electron group

geometry around each central atom. 3. Select the hybridization scheme that matches the

electron group geometry. 4. Sketch the atomic and hybrid orbitals on the atoms

in the molecule, showing overlap of the appropriate orbitals

5. Label the bonds as σ or π

Predict the hybridization and bonding scheme for CH2CH2

1. Start by drawing the Lewis structure

2. Use VSEPR Theory to predict the electron group geometry around each central atom

The molecule has two interior atoms. Since each atom has three electron groups (one double bond and two single bonds), the electron geometry about each atom is trigonal planar.

Predict the hybridization and bonding scheme for CH2CH2

3. Select the hybridization scheme that matches the electron group geometry

4. Sketch the atomic and hybrid orbitals on the atoms in the molecule, showing overlap of the appropriate orbitals

C1 = trigonal planar C1 = sp2

C2 = trigonal planar C2 = sp2

continued…

Predict the hybridization and bonding scheme for CH3CHO

C1 = 4 electron areas C1= tetrahedral C2 = 3 electron areas C2 = trigonal planar

1. Start by drawing the Lewis structure

2. Use VSEPR Theory to predict the electron group geometry around each central atom

Predict the hybridization and bonding scheme for CH3CHO

3. Select the hybridization scheme that matches the electron group geometry

4. Sketch the atomic and hybrid orbitals on the atoms in the molecule, showing overlap of the appropriate orbitals

C1 = tetrahedral C1 = sp3

C2 = trigonal planar C2 = sp2

Label the bonds as σ or π

Predict the hybridization and bonding scheme for CH3CHO

σ

HC C

OH

HH

π

Additional Examples Follow

Predict the hybridization of all the atoms in H3BO3

H = can’t hybridizeB = 3 electron groups = sp2

O = 4 electron groups = sp3

Predict the hybridization and bonding scheme of all the atoms in NClO

• •

O N C l •

•

•

•

• • • • • •

N = 3 electron groups = sp2

O = 3 electron groups = sp2

Cl = 4 electron groups = sp3

Cl

O N

Cl

O N

SOF4 Orbital Diagram

↑ sp3d S ↑ ↑

↑ ↑

sp2 O ↑ ↑↓ ↑

2p F

σ

σ

d S p O

π

↑ ↑ ↑↓ ↑

2p F

σ

↑

2p F

σ

↑

2p F

σ

sp3d hybridization

Atom with five electron groups around ittrigonal bipyramid electron geometry Seesaw, T–Shape, Linear 120° & 90° bond angles

Use empty d orbitals from valence shelld orbitals can be used to make π bonds