Notes Lecture 01 Ch 14

8/6/2019 Notes Lecture 01 Ch 14

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Chemistry 162

Text: Chemical Principles, 6th Ed.- By Steven Zumdahl

Chapter #14 : Covalent Bonding: Orbitals

Chapter #15 : Chemical Kinetics

Chapter #16 : Liquids and Solids

Chapter #17 : Properties of Solutions

Chapters #18: Representative Elements

Chapter #19: Transition Metals and Coordination Chemistry

Chapter #14 - Covalent Bonding: Orbitals

14.1) Hybridization and the Localized Electron Model

14.2) The Molecular Orbital Model

14.3) Bonding in Homonuclear Diatomic Molecules

14.4) Bonding in Heteronuclear Diatomic Molecules

14.5) Combining the Localized Electron and Molecular Orbital

Models

14.6) Orbitals: Human Inventions

14.7) Molecular Spectroscopy



Hybridization and the Localized Electron Model

Chapter 12:Atomic orbitals, Properties of electrons, Wave functions,

Electronic configurations, Aufbau principle, etc.

Chapter 13:

General Concepts of Bonding in Molecules- Types of bonds: ionic, covalent, etc.

- Bond energies, lengths, polarities, etc.

Localized Electron Model

- Lewis dot structures- Resonance structures

- The octet rule

- VSEPR model

Chapter 14:

Central Themes of Valence Bond Theory

(Localized Electron Model)

1) Maximum overlap. The bond strength depends on the coulombic

attraction between the shared electrons and the two nuclei. The

greater the orbital overlap, the stronger the bond.

Basic Principle of Valence Bond Theory: A covalent bond formswhen the orbitals from two atoms overlap and a pair of electronsoccupies the region between the two nuclei.



Central Themes of Valence Bond Theory

(Localized Electron Model)

1) Maximum overlap. The bond strength depends on the coulombic

attraction between the shared electrons and the two nuclei. The

greater the orbital overlap, the stronger the bond.

2) Spins pair. The two electrons in the overlap region occupy thesame space and therefore must have opposite spins. (Pauli exclusion

principle)

Basic Principle of Valence Bond Theory: A covalent bond formswhen the orbitals from two atoms overlap and a pair of electronsoccupies the region between the two nuclei.

3) Hybridization. To explain experimental observations, Pauling proposed that the valence atomic orbitals in a molecule are

different from those in the isolated atoms. We call thisHybridization!



Example: Methane

• 4 equivalent C-H covalent bonds• VSEPR predicts a tetrahedral geometry

Lewis dot structure VSEPR 3D shape

Recall, the method of generating

Lewis dot structures

total # valence

electrons = 8

0 e – left

H CH HH

H C

H

H

H

CH4

1. determine the total # of valence electrons

2. use pairs of e – to connect terminal atoms tocentral atom, and then distribute the rest as

lone pairs



Recall, the method of determining

3–D structure (VSEPR)

9

• Valence Shell Electron Pair Repulsion (VSEPR):

the!3-D geometric structure is determined by

minimizing repulsion of electron pairs (both bond

pairs and lone pairs).

H C

H

H

H

Lewis Structure VSEPR Structure

(VSEPR)



(VSEPR)

Carbon: 2s22p2

So, how do we explain the formation of

4 equivalent C-H bonds?

The Valence Orbitals of a Carbon Atom



Hybridization: Mixing of Atomic Orbitals to

form New Orbitals for Bonding

4 atomic orbitals 4 new hybrid orbitals

Grouped together, the four have

a tetrahedral shape



!1

= 1/2[(2s) + (2px) + (2p

y) + (2p

z)]

!2

= 1/2[(2s) + (2px) - (2p

y) - (2p

z)]

!3

= 1/2[(2s) - (2px) + (2p

y) - (2p

z)]

!4

= 1/2[(2s) - (2px) - (2p

y) + (2p

z)]

Other Representations of Hybridization:

Hybridization is related to the number of

electron pairs determined from VSEPR:Methane

VSEPR:

tetrahedral

sp3 hybridized

Ammonia

VSEPR:

tetrahedralsp3 hybridized

Water

VSEPR :

tetrahedral

sp3 hybridized





used for bonding

remaining p orbital used for bonding

Problem: Describe the hybridization and bonding of the

carbon orbitals in ethylene (C2H4)

VSEPR:

trigonal planarC C

H

H

H

H

The Formation of Involves the

Combination of an s–orbital with a px – and py –orbital

three atomic orbitals

three hybrid

orbitals



One of the p–orbitals, the pz –orbital remains

unhybridized

The unhybridized pz –orbital will be used

to form a ! –bond

z

The ! (sigma) bonds in C2H4



A carbon-carbon double bond consists

of a ! bond and a " bond

C C

H

H

H

H

bonding vs bondingRead pages 665 of Zumdahl

• Two modes of bonding are important for 1st and 2nd row elements: # bonding and " bonding

• These two differ in their relationship to the internuclear • axis:

# bonds have electron density ON the internuclear axis

" bonds have electron density ABOVE AND BELOW the internuclear axis

" bond



Bonding in ethylene (C2H4)

Problem: Describe the hybridization and bonding of the

carbon orbitals in Carbon Dioxide (CO2)

VSEPR:

linear

sp hybridized orbitals for # bonding



The ! (sigma) bonds in CO2

Two of the p–orbitals of CO2 remain

unhybridized

The unhybridized orbitals will be used

to form two ! –bonds



Bonding in Carbon Dioxide (CO2)

5- and 6-Coordinate Molecules: e.g. PCl5

VSEPR: AB5

trigonal bipyrimidal

Expanded Octet for P

(10 valence electrons)



1. Molecular Formula

2. Lewis Structure

3. VSEPR- shape and arrangement

4. Hybrid orbitals

Strategies used in forming

Hybrid Orbitals

Summary:hybrid

combination

sp

sp2

sp3



Summary:

hybrid

combination

dsp3

d2sp3

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Notes Lecture 01 Ch 14