AP Chem Unit 9: Covalent Bonding Orbitals
Guided Notes
Main Ideas and Sections:
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Hybridization and the Localized Electron Model
The Molecular Orbital Model
Bonding in Homonuclear Diatomic Molecules
Bonding in Heteronuclear Diatomic Molecules
Combining the Localized Electron and MO models
Hybridization and the Localized Electron Model
Localized Electron Model
The arrangement of valence electrons is represented by the Lewis structure or
structures, and the molecular geometry can be predicted from the VSEPR model.
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Atomic orbitals are used to share electrons and form bonds
Hybridization
In general we assume that bonding involves only the valence orbitals.
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The mixing of atomic orbitals to form special bonding orbitals is called
hybridization.
o Carbon is said to undergo sp3 hybridization or is sp3 hybridized
because is uses one s orbital and three p orbitals to form four identical
bonding orbitals.
o 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 arrangement.
sp3 hybridized
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sp2 Hybridization
Example: Ethylene (C2H4) is commonly used in plastics and has a C=C double
bond. Each carbon uses sp2 hybridization in this molecule because a double
bond acts as one effective pair.
o 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 orbitals.
o The double bond utilizes one sigma bond that is hybridized and one pi
bond with the unhybridized p orbital.
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Example: sp2 Hybridization: Ethylene
Multiple Bonds
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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 unhybridized orbitals.
Example: sp2 Hybridization: Ethylene
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sp Hybridization
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sp hybridization involves one s orbital and one p orbital. Two effective pairs
will always require sp hybridization.
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CO2 is sp hybridized
Example: sp Hybridization: CO2
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Example: sp Hybridization: N2
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sp3d Hybridization
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When a molecule exceeds the octet rule, hybridization occurs using d
orbitals. Also called dsp3.
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PCl5 has sp3d hybridization and is trigonal bipyramidal.
Example: sp3d Hybridization: PCl5
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Each chlorine atom displays a tetrahedral arrangement around the atom.
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sp3d2 Hybridization
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An octahedral arrangement requires six effective pairs around the central
atom.
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SF6 has sp3d2 hybridization.
Example Problem: How is the xenon atom in XeF4 hybridized?
Localized Electron Summary
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Draw the Lewis Structure
Determine the arrangement of electron pairs using the VSEPR model.
Specify the hybrid orbitals needed to accommodate the electron pairs.
o 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.
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Effective Pairs and Their Spatial Arrangement
The Molecular Orbital Model
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.
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Since we do not know the details of the electron movements, we cannot deal
with the electron-electron interactions in a specific way
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The Molecular Orbital model helps us to deal with the molecular problem.
Molecular Orbitals
Molecular orbitals (MOs) have many of the same characteristics as atomic orbitals.
Two of the most important are:
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MOs can hold two electrons with opposite spins.
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The square of the MO’s wave function indicates electron probability.
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For simplicity we will first look at the H2 molecule.
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The combination of hydrogen 1s atomic orbitals results in 2 molecular
orbitals.
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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.
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When the orbitals are added, the matching phases produce constructive
interference and the opposite phases produce destructive interference.
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A constructive combination gives a bonding MO. This gives an enhanced
electron probability between the nuclei.
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The destructive combination gives an antibonding MO. This interference
produces a node between the nuclei.
Two MOs exist for H2:
o MO1= 1sH1 + 1sH2
 MO1 is constructive and therefore a bonding MO
 MO1 is lower energy
o MO2 = 1sH1 – 1sH2
 MO2 is destructive and therefore an antibonding MO
 MO2 is higher energy
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MOs
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.
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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 orbitals.
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 probonding.
If two electrons were forced to occupy the higher-energy MO2 this would be antibonding and the lower energy of the separated atoms would be favored.
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Bonding and Antibonding
Labels are given to MOs indicate their symmetry (shape), the parent atomic orbitals,
and whether they are bonding or antibonding.
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Antibonding character is indicated by an asterisk.
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Subscripts indicate parent orbitals
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σ and π indicate shape.
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H2 has the following MOs:
o MO1 = σ1s
o 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: σ1s2
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Each molecular orbital can hold two electrons, but the spins must be
opposite.
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Orbitals are conserved. The number of molecular orbitals will always be the
same as the number of atomic orbitals used to construct them.
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From this molecular electron configuration, we can determine a molecules stability.
Example: Would H2- be stable?
o (σ1s )2 ( σ1s*) 1
o 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- is formed.
o 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.
Bond Order
To indicate bond strength, we use the concept of bond order.
Example: H2 has a bond order of 1
H2- has a bond order of ½
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.
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Larger bond order means greater bond strength.
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Bond order of 0 gives us a molecule that doesn’t exist.
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Bonding in Homonuclear Diatomic Molecules
Homonuclear Bonding
When looking at bonding beyond energy level 1, we need to consider what orbitals
are overlapping and therefore bonding.
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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.
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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.
Example: What is the molecular electron configuration and bond order of Li2?
Example: What is the molecular electron configuration and bond order of
Be2?
MOs from p orbitals
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p orbitals must overlap in such a way that the wave patterns produce
constructive interference. As with the s orbitals, the destructive interference
produces a node in the wave pattern and decreases the probability of
bonding.
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When the parallel p orbitals are combined with the positive and negative
phases matched, constructive interference occurs, giving a bonding π orbital.
When orbitals have opposite phases, destructive interference results in an
antibonding π orbital.
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MOs from p orbitals continued
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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.
Example: B2
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Exceptions
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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 electrons.
Paramagnetism
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:
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Paramagnetism – causes the substance to be attracted into the magnetic field.
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Diamagnetism – causes the substance to be repelled from the magnetic field.
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Paramagnetism
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Paramagnetism is associated with unpaired electrons and diamagnetism is
associated with paired electrons.
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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 diamagnetism.
Summary
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There are definite correlations between bond order, bond energy, and bond
length. As bond order increases so does bond energy and bond length
decreases.
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Comparison of bond orders between different molecules cannot predict bond
energies of different molecules.
o B2 and F2 both have bond order of 1 but bond energies are very
different. B-B is a much stronger bond.
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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.
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Example Problem: For O2, O2+, and O2-, give the MO electron configuration
and the bond order for each. Which has the strongest bond?
MO order with mixing
MO order without mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
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Example Problem: Use the molecular orbital model to predict the bond order
and magnetism of each of the following molecules: Ne2 and P2
MO order with mixing
MO order without mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
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Bonding in Heteronuclear Diatomic Molecules
Heteronuclear Molecules
When dealing with different atoms within diatomic molecules we can still use the
MO model to determine bond order and magnetism
Example: NO
Example Problem: Use the MO model to predict the magnetism and bond
order of the NO+ and CN- ions.
Heteronuclear Diatomics
What happens with the diatomic molecules are very different?
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A molecular orbital forms between two different atomic orbitals.
o HF example
o Note: energy level difference vs. electronegativity
o
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Combining the Localized Electron and MO models
Resonance
When a molecule has resonance. It is usually a double bond that can have different
positions around the molecule.
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The single σ bonds remain localized and the π bonds are said to be
delocalized.
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Benzene: All C-C bonds are known to be equivalent and the molecule has
resonance
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Benzene:
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Resonance continued
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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.
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This is called delocalized π bonding.
NO3NO3- ion also displays delocalized π bonding
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MO order with mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order with mixing
σ2p*
π2p*
σ2p
π2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
MO order without mixing
σ2p*
π2p*
π2p
σ2p
σ2s*
σ2s
σ1s*
σ1s
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