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Lecture 1
Chemical Bonds: Atomic Orbital
Theory and Molecular Orbital
Theory
Dr. A.K.M. Shafiqul Islam
14.07.08
Summary of Modern Atomic Theory
 Atoms have an internal structure consisting of one
or more subatomic particles: protons, neutrons,
and electrons.
proton  positive charge
mass = 1.673 x 10-27 kg
neutron  no charge
mass = 1.675 x 10-27 kg
electron  negative charge
mass = 9.109 x 10-31 kg
Summary of Modern Atomic Theory
 Most of the mass of an atom is concentrated in the
nucleus.
 The nucleus contains one or more positively
charged protons, and one or more neutrons with
no electrical charge.
Summary of Modern Atomic Theory
 One or more negatively
charged electrons are in
constant motion
somewhere outside the
nucleus.
 The number of electrons
is equal to the number of
protons; the atom has no
overall electrical charge.
Summary of Modern Atomic Theory
 An atom is mostly free
space because the volume
of the nucleus and the
electrons outside the
nucleus are extremely
small compared to the
overall volume of the
atom.
Summary of Modern Atomic Theory
 Heisenberg Uncertainty Principle
• It is not possible to determine both the position and the
momentum of an electron.
 Wave Function
y
• Describes the energy of an electron and the probability
of finding the electron in a region around the nucleus.
 Atomic Orbital
• The probability distribution about one atomic nucleus
(i.e., wave function = atomic orbital).
– Amplitudes – numerical magnitudes
– Signs – positive or negative
– Nodes – values of wavefunction equals zero (given by
quantum numbers)
Summary of Modern Atomic Theory
 Heisenberg Uncertainty Principle
• It is not possible to determine both the position and the
momentum of an electron.
 Wave Function
y
• Describes the energy of an electron and the probability
of finding the electron in a region around the nucleus.
 Atomic Orbital
• The probability distribution about one atomic nucleus
(i.e., wave function = atomic orbital).
– Amplitudes – numerical magnitudes
– Signs – positive or negative
– Nodes – values of wavefunction equals zero (given by
quantum numbers)
Summary of Modern Atomic Theory
 Heisenberg Uncertainty Principle
• It is not possible to determine both the position and the
momentum of an electron.
 Wave Function
y
• Describes the energy of an electron and the probability
of finding the electron in a region around the nucleus.
 Atomic Orbital
• The probability distribution about one atomic nucleus
(i.e., wave function = atomic orbital).
– Amplitudes – numerical magnitudes
– Signs – positive or negative
– Nodes – values of wavefunction equals zero (given by
quantum numbers)
Summary of Modern Atomic Theory
 Heisenberg Uncertainty Principle
• It is not possible to determine both the position and the
momentum of an electron.
 Wave Function
y
• Describes the energy of an electron and the probability
of finding the electron in a region around the nucleus.
 Atomic Orbital
• The probability distribution about one atomic nucleus
(i.e., wave function = atomic orbital).
– Amplitudes – numerical magnitudes
– Signs – positive or negative
– Nodes – values of wave function equals zero (given
by quantum numbers)
Atomic Orbitals
s orbital
p orbitals
px
py
pz
Hybrid Atomic Orbitals
hybrid
sp3()
sp2 ()
sp ()
bonding
head-head
sideways
sideways
shape
rotation bond
tetrahedral free
single
triangular rigid
double
linear
rigid
triple
Carbon
Atomic Number = 6
2p
2s
1s
Carbon
Atomic Number = 6
2p
2s
1s
Oxygen
Atomic Number = 8
2p
2s
1s
Oxygen
Atomic Number = 8
2p
2s
1s
Chemical Bonding
Chemical Bonds
• The forces holding atoms together in
compounds.
Valence Electrons
• The electrons in the outer shell.
Lewis Dot Representation of Atoms
• Dots around the chemical symbol of an atom
represent the valence electrons.
Chemical Bonding
Chemical Bonds
• The forces holding atoms together in
compounds.
Valence Electrons
• The electrons in the outer shell.
Lewis Dot Representation of Atoms
• Dots around the chemical symbol of an atom
represent the valence electrons.
Chemical Bonding
Chemical Bonds
• The forces holding atoms together in
compounds.
Valence Electrons
• The electrons in the outer shell.
Lewis Dot Representation of Atoms
• Dots around the chemical symbol of an atom
represent the valence electrons.
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Examples
Atom
Electronic
Structure
Electronic
Configuration
Lewis Dot
Structure
3p
3s
Boron
2
2
1
1s 2s 2p
2p
2s
B
1s
3p
3s
Phosphorus
[Ne] 3s23p3
2p
2s
1s
P
Molecular Orbital Theory
 Molecular orbitals – delocalized over entire
molecule.
 First Principle
• The total number of molecular orbitals is always
equal to the total number of atomic orbitals
contributed by the atoms that have combined.
Molecular Orbitals for H2
 Bonding molecular orbital
• addition of two atomic orbitals, leads to the probability of finding an
e- between the atoms.
 Antibonding molecular orbital
• substration of one atomic orbital from the other, leads to reduced
probability of finding an e- between the nuclei, increased in other
regions.
Molecular Orbital Theory
 Second Principle
• The bonding molecular orbital is lower in energy
than the parent orbitals, and the antibonding
orbital is higher in energy.
 Third Principle
• The electrons of the molecule are assigned to
orbitals of successively higher energy
(according to Pauli exclusion principle and the
Hund’s rule).
The Ionic Bond
Transfer of Electrons from One Atom to
Another
Example:
• Sodium chloride
Na
Cl + 1 e
Na
+ 1e
Cl
The Ionic Bond
The electrostatic attraction between
oppositely charged ions.
The Covalent Bond
Some atoms do not transfer electrons from
one atom to another to form ions.
Instead they form a chemical bond by
sharing pairs of electrons between them.
A covalent bond consists of a pair of
electrons shared between two atoms.
The Covalent Bond
Some atoms do not transfer electrons from
one atom to another to form ions.
Instead they form a chemical bond by
sharing pairs of electrons between them.
A covalent bond consists of a pair of
electrons shared between two atoms.
The Covalent Bond
Some atoms do not transfer electrons from
one atom to another to form ions.
Instead they form a chemical bond by
sharing pairs of electrons between them.
A covalent bond consists of a pair of
electrons shared between two atoms.
Hydrogen, H2
*
E*
1s
1s
E

Fluorine, F2
Fluorine, F2
F + F
F
F
Fluorine, F2
F + F
F
F
F2 *
F 2p
F 2p
F2 
Nitrogen, N2
Nitrogen, N2
N
O +
+
C +
N N
N
O
O C O
Nitrogen, N2
N
+
N N
N
Carbon Dioxide, CO2
O +
C +
O
O C O
Nitrogen, N2
N
+
N N
N
Carbon Dioxide, CO2
O +
C +
O
O C O
Molecular Orbitals for He2
He2 has no net stability – two He atoms
have no tendency to combine.