3. Localized Electron Model
The arrangement of valence electrons is represented by
the Lewis structure or structures, and the molecular
geometry can be predicted from theVSEPR model.
Atomic orbitals are used to share electrons and form
In general we assume that bonding involves only the valence
The mixing of atomic orbitals to form special bonding orbitals
is called hybridization.
Carbon is said to undergo sp3
hybridization or is sp3
because is uses one s orbital and three p orbitals to form four
identical bonding orbitals.
The four sp3
orbitals are identical in shape each one having a
large lobe and a small lobe. The four orbitals are oriented in
space so that the large lobes form a tetrahedral
Ethylene (C2H4) is commonly used in plastics and has a
C=C double bond. Each carbon uses sp2
this molecule because a double bond acts as one
In forming the sp2
orbitals, one 2p orbital on carbon has
not been used. This remaining p orbital is oriented
perpendicular to the plane of the sp2
The double bond utilizes one sigma bond that is
hybridized and one pi bond with the unhybridized p
14. Multiple Bonds
Single bonds are sigma bonds (σ) and the electron pair is
shared in an area centered on a line running between
the atoms.These are hybridized bonding orbitals.
With multiple bonds, a sigma bond is formed and then
one or two pi bond (π) form. These electrons occupy the
space above and below the sigma bond and use
26. Localized Electron Summary
Draw the Lewis Structure
Determine the arrangement of electron pairs using the
Specify the hybrid orbitals needed to accommodate the
Do not overemphasize the characteristics of the
separate atoms. It is not where the valence electrons
originate that is important; it is where they are
needed in the molecule to achieve stability.
29. Molecular Orbital Model
The localized electron model works very well with the
prediction of structure and bonding of molecules, but
the electron correlation problem still exists.
Since we do not know the details of the electron
movements, we cannot deal with the electron-electron
interactions in a specific way
The Molecular Orbital model helps us to deal with the
30. Molecular Orbitals
Molecular orbitals (MOs) have many of the same
characteristics as atomic orbitals.Two of the most
MOs can hold two electrons with opposite spins.
The square of the MO’s wave function indicates electron
For simplicity we will first look at the H2 molecule.
The combination of hydrogen 1s atomic orbitals results in 2
The wave phases of the atomic orbitals combine/overlap.
Since electrons move in wave functions, this causes
constructive and destructive interference in the wave pattern.
When the orbitals are added, the matching phases produce
constructive interference and the opposite phases produce
A constructive combination gives a bonding MO. This
gives an enhanced electron probability between the
The destructive combination gives an antibonding MO.
This interference produces a node between the nuclei.
Two MOs exist for H2:
MO1= 1sH1 + 1sH2
MO1 is constructive and therefore a bonding MO
MO1 is lower energy
MO2 = 1sH1 – 1sH2
MO2 is destructive and therefore an antibonding MO
MO2 is higher energy
The type of electron distribution described in these MOs is
called sigma as in the localized electron model. MO1 and
MO2 are sigma (σ) molecular orbitals.
In this molecule only the molecular orbitals are available
for occupation by electrons. The 1s atomic orbitals of
the hydrogen atoms no longer exist, because the H2
molecule – a new entity – has its own set of new
The energy level of the bonding
MO is lower and more stable than
that of the antibonding MO. Since
molecule formation favors the
lowest energy state, this provides
the driving force for molecule
formation of H2. This is called
If two electrons were forced to
occupy the higher-energy MO2 this
would be anti-bonding and the
lower energy of the separated
atoms would be favored.
Labels are given to MOs indicate their symmetry (shape), the
parent atomic orbitals, and whether they are bonding or
Antibonding character is indicated by an asterisk.
Subscripts indicate parent orbitals
σ and π indicate shape.
H2 has the following MOs:
MO1 = σ1s
MO2 = σ1s*
Molecular electron configurations can be written in much
the same way as atomic (electron) configurations. Since
the H2 molecule has two electrons in the σ1s molecular
orbital, the electron configuration is: σ1s
Each molecular orbital can hold two electrons, but the
spins must be opposite.
Orbitals are conserved. The number of molecular
orbitals will always be the same as the number of
atomic orbitals used to construct them.
From this molecular electron configuration, we can determine a
The key idea is that H2
would exist if it were a lower energy than its
separated parts. Two electrons are in bonding and one is in
antibonding. Since more electrons favor bonding H2
This also is a good indicator of bond strength. H2 has a stronger
bond than H2
. The net lowering of the bonding electrons by one is a
direct relationship to bond strength. H2 is twice as strong.
40. Bond Order
To indicate bond strength, we use the concept of bond
Example: H2 has a bond order of 1
has a bond order of ½
Bond Order =
− #οφαντιβονδινγ ε−
41. Bond Order
Bond order is an indication of bond strength because it
reflects the difference between the number of bonding
electrons and the number of antibonding electrons.
Larger bond order means greater bond strength.
Bond order of 0 gives us a molecule that doesn’t exist.
43. Homonuclear Bonding
When looking at bonding beyond energy level 1, we need to
consider what orbitals are overlapping and therefore
Li2 has electrons in the 1s and 2s orbitals; the 2s orbitals
are much larger and overlap, but the 1s orbitals are
smaller and do not overlap.
To participate in molecular orbitals, atomic orbitals must
overlap in space. This means that only the valence
orbitals of the atoms contribute significantly to the
molecular orbitals of a particular molecule.
44. Li2 MO
What is the molecular electron configuration and bond
order of Li2?
with a bond order of 1
Li2 is a stable molecule because the overall energy of the
molecule is lower than the separate atoms.
45. Be2 MO
What is the molecular electron configuration and bond
order of Be2?
with a bond order of 0
Be2 has 2 bonding and 2 antibonding electrons and is not
more stable than the individual atoms. Be2 does not
46. MOs from p orbitals
p orbitals must overlap in
such a way that the wave
As with the s orbitals, the
produces a node in the
wave pattern and
decreases the probability
47. MOs from p orbitals
When the parallel p
orbitals are combined with
the positive and negative
occurs, giving a bonding π
orbital. When orbitals
have opposite phases,
results in an antibonding π
48. MOs from p orbitals
Since the electron probability
lies above and below the line
between the nuclei (with
parallel p orbitals), the
stability of a π molecular
bonding orbital is less than
that of a σ bonding orbital.
Also, the antibonding π MO is
not as unstable as the
antibonding σ MO. The
energies associated with the
orbitals reflect this stability.
B2, C2, and N2 molecules use the
same set of molecular orbitals
that we expect but some
mixing of orbital energies
occurs. The s and p atomic
orbitals mix or hybridize in a
way that changes some MO
energy states. This affects
filling order and pairing of
Most materials have no magnetism until they are placed in
a magnetic field. However, in the presence of such a
field, magnetism of two types can be induced:
Paramagnetism – causes the substance to be attracted
into the magnetic field.
Diamagnetism – causes the substance to be repelled
from the magnetic field.
Paramagnetism is associated with unpaired electrons
and diamagnetism is associated with paired electrons.
Any substance that has both paired and unpaired
electrons will exhibit a net paramagnetism since the
effect of paramagnetism is much stronger than that of
There are definite correlations between bond order, bond
energy, and bond length. As bond order increases so does
bond energy and bond length decreases.
Comparison of bond orders between different molecules
cannot predict bond energies of different molecules.
B2 and F2 both have bond order of 1 but bond energies are
very different. B-B is a much stronger bond.
N2 has a bond order of 3 and has a very large bond energy. N2 is
a very stable molecule and is used to drive powerful reactions.
56. Example Problem
For O2, O2
, and O2
, give the MO
electron configuration and the
bond order for each. Which has
the strongest bond?
O2 : (σ2s)2
B2 C2 N2 O2 F2
57. Example Problem
Use the molecular orbital model to predict the bond order
and magnetism of each of the following molecules: Ne2
Ne2 bond order is 0: does not exist
P2 bond order is 3 and diamagnetic
B2 C2 N2 O2 F2
59. Heteronuclear Molecules
When dealing with different atoms within diatomic
molecules we can still use the MO model to determine
bond order and magnetism
60. NO example
The valence electrons from
both atoms fill in the order
expected by the model.
The bond order is 2.5 and is
61. Example Problem
Use the MO model to predict the
magnetism and bond order of
Both ions are diamagnetic and
have the same configuration.
Their bond order is 3
62. Heteronuclear Diatomics
What happens with the
diatomic molecules are
A molecular orbital forms
between two different
Note: energy level
When a molecule has resonance. It is usually a double
bond that can have different positions around the
The single σ bonds remain localized and the π bonds are
said to be delocalized.
Benzene: The p orbitals are perpendicular to the plane and
form π bonds above and below the plane.The electrons in the
π bonds delocalize and give six equivalent C-C bonds that give
the structure true resonance.
This is called delocalized π bonding.