Covalent Bonding Orbitals Adapted from bobcatchemistry

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Covalent Bonding OrbitalsAdapted from
bobcatchemistry
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Hybridization and the Localized Electron Model
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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.
• Atomic orbitals are used to share electrons and
  form bonds

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Hybridization

• In general we assume that bonding involves only
  the valence orbitals.
• The mixing of atomic orbitals to form special
  bonding orbitals is called hybridization.
• 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.
• 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.

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Hybridization
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Hybridization
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sp3 example Methane
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sp3 example Ammonia
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sp2 Hybridization
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sp2 Hybridization
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sp2 Hybridization

• Ethylene (C2H4) is commonly used in plastics and
  has a CC double bond. Each carbon uses sp2
  hybridization in this molecule because a double
  bond acts as one effective pair.
• 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.
• The double bond utilizes one sigma bond that is
  hybridized and one pi bond with the unhybridized
  p orbital.

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sp2 Hybridization
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sp2 Hybridization Ethylene
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Multiple Bonds

• Single bonds are sigma bonds (s) 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 (p) form. These
  electrons occupy the space above and below the
  sigma bond and use unhybridized orbitals.

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sp2 Hybridization Ethylene
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sp Hybridization

• sp hybridization involves one s orbital and one p
  orbital. Two effective pairs will always require
  sp hybridization.
• CO2 is sp hybridized

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sp Hybridization
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sp Hybridization CO2
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sp Hybridization CO2
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sp Hybridization CO2
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sp Hybridization N2
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sp3d Hybridization

• When a molecule exceeds the octet rule,
  hybridization occurs using d orbitals. Also
  called dsp3.
• PCl5 has sp3d hybridization and is trigonal
  bipyramidal.

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sp3d Hybridization PCl5

• Each chlorine atom displays a tetrahedral
  arrangement around the atom.

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sp3d2 Hybridization

• An octahedral arrangement requires six effective
  pairs around the central atom.
• SF6 has sp3d2 hybridization.

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Example Problem

• How is the xenon atom in XeF4 hybridized?

sp3d2
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Localized Electron Summary

• Draw the Lewis Structure
• Determine the arrangement of electron pairs using
  the VSEPR model.
• Specify the hybrid orbitals needed to accommodate
  the electron pairs.
• 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
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The Molecular Orbital Model
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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 molecular problem.

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Molecular Orbitals

• Molecular orbitals (MOs) have many of the same
  characteristics as atomic orbitals. Two of the
  most important are
• MOs can hold two electrons with opposite spins.
• The square of the MOs wave function indicates
  electron probability.

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MOs

• For simplicity we will first look at the H2
  molecule.
• The combination of hydrogen 1s atomic orbitals
  results in 2 molecular orbitals.
• 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 destructive interference.
  

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MOs

• A constructive combination gives a bonding MO.
  This gives an enhanced electron probability
  between the nuclei.
• The destructive combination gives an antibonding
  MO. This interference produces a node between
  the nuclei.

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MOs

• 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

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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 (s)
  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 orbitals.

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MOs

• 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 anti-bonding and
  the lower energy of the separated atoms would be
  favored.

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Bonding and Antibonding
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MOs

• Labels are given to MOs indicate their symmetry
  (shape), the parent atomic orbitals, and whether
  they are bonding or antibonding.
• Antibonding character is indicated by an
  asterisk.
• Subscripts indicate parent orbitals
• s and p indicate shape.
• H2 has the following MOs
• MO1 s1s
• MO2 s1s

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MOs

• Molecular electron configurations can be written
  in much the same way as atomic (electron)
  configurations. Since the H2 molecule has two
  electrons in the s1s molecular orbital, the
  electron configuration is s1s2
• 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.

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MOs

• From this molecular electron configuration, we
  can determine a molecules stability.
• Would H2- be stable?
• (s1s )2 ( s1s) 1
• 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.
• 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.

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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

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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 doesnt
  exist.

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Bonding in Homonuclear Diatomic Molecules
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Homonuclear Bonding

• When looking at bonding beyond energy level 1, we
  need to consider what orbitals are overlapping
  and therefore bonding.
• 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.

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Li2 MO

• What is the molecular electron configuration and
  bond order of Li2?
• s2s2 with a bond order of 1
• Li2 is a stable molecule because the overall
  energy of the molecule is lower than the separate
  atoms.

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Be2 MO

• What is the molecular electron configuration and
  bond order of Be2?
• (s2s)2 (s2s)2 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 form.

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MOs from p orbitals

• 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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MOs from p orbitals

• When the parallel p orbitals are combined with
  the positive and negative phases matched,
  constructive interference occurs, giving a
  bonding p orbital. When orbitals have opposite
  phases, destructive interference results in an
  antibonding p orbital.

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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 p molecular
  bonding orbital is less than that of a s bonding
  orbital. Also, the antibonding p MO is not as
  unstable as the antibonding s MO. The energies
  associated with the orbitals reflect this
  stability.

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B2 Example

• 1s2 2s2 2p1
• 1s2 does not bond
• 2s2 and 2p1 bond
• (s2s)2 (s2s)2 (s2p)2
• Bond order 1

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Exceptions

• 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.

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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
• Paramagnetism causes the substance to be
  attracted into the magnetic field.
• Diamagnetism causes the substance to be
  repelled from the magnetic field.

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Paramagnetism
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Paramagnetism

• 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 diamagnetism.
• Paramagnetism Video

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Summary

• 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.

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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?
• O2 (s2s)2 (s2s)2 (s2p)2 (p2p)4 (p2p)2
  BO2
• O2 (s2s)2 (s2s)2 (s2p)2 (p2p)4 (p2p)1
  BO2.5
• O2- (s2s)2 (s2s)2 (s2p)2 (p2p)4 (p2p)3
  BO1.5

B2 C2 N2 B2 C2 N2 B2 C2 N2 O2 F2 O2 F2 O2 F2
s2p   s2p  
p2p   p2p  
s2p   p2p  
p2p   s2p  
s2s   s2s  
s2s   s2s  
s1s   s1s  
s1s     s1s    
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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
• Ne2 bond order is 0 does not exist
• P2 bond order is 3 and diamagnetic

B2 C2 N2 B2 C2 N2 B2 C2 N2 O2 F2 O2 F2 O2 F2
s2p   s2p  
p2p   p2p  
s2p   p2p  
p2p   s2p  
s2s   s2s  
s2s   s2s  
s1s   s1s  
s1s     s1s    
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Bonding in Heteronuclear Diatomic Molecules
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Heteronuclear Molecules

• When dealing with different atoms within diatomic
  molecules we can still use the MO model to
  determine bond order and magnetism

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NO example

• The valence electrons from both atoms fill in the
  order expected by the model.
• The bond order is 2.5 and is paramagnetic.

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Example Problem

• Use the MO model to predict the magnetism and
  bond order of the NO and CN- ions.
• Both ions are diamagnetic and have the same
  configuration. Their bond order is 3

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Heteronuclear Diatomics

• What happens with the diatomic molecules are very
  different?
• A molecular orbital forms between two different
  atomic orbitals.
• HF example
• Note energy level difference vs.
  electronegativity

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Combining the Localized Electron and MO models
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Resonance

• When a molecule has resonance. It is usually a
  double bond that can have different positions
  around the molecule.
• The single s bonds remain localized and the p
  bonds are said to be delocalized.

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Resonance

• Benzene All C-C bonds are known to be
  equivalent and the molecule has resonance

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Resonance

• Benzene The s bonds remain centered (on the
  plane) between C atoms.

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Resonance

• Benzene The p orbitals are perpendicular to the
  plane and form p bonds above and below the plane.
  The electrons in the p bonds delocalize and give
  six equivalent C-C bonds that give the structure
  true resonance.
• This is called delocalized p bonding.

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NO3-

• NO3- ion also displays delocalized p bonding.

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The End
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MO Worksheet

B2 C2 N2 B2 C2 N2 B2 C2 N2 O2 F2 O2 F2 O2 F2
s2p   s2p  
p2p   p2p  
s2p   p2p  
p2p   s2p  
s2s   s2s  
s2s   s2s  
s1s   s1s  
s1s     s1s