CarbonCarbon bonds: cis and trans isomers

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Chemical Bonding Theory

• Carbon-Carbon bonds cis- and trans- isomers
• Rotation about the internuclear axis of C-C
  single bonded atoms is relatively easy at room
  temperature
• Rotation about the internuclear axis of CC
  double bonded atoms is not very easy at room
  temperature p-bond energy 260 kJ/mol
• The result can be that two structural forms are
  frozen out
• Each form is actually a different compound
  because the atoms have different arrangements in
  3-space
• Each can have different physical properties such
  as boiling temperature, molecular polarity or
  color
• Isomers are compounds that have the same
  molecular formula but different arrangements of
  the atoms in 3-space
• Example cis- and trans- 1,2-dichloroethene

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Chemical Bonding Theory

• Delocalized p bonds When two or more resonance
  structures can be written involving p bonds, the
  p bonds can be represented as being spread out
  over the molecule where the p bonds occur.
• The electrons and the bonds containing them are
  delocalized.
• Example Benzene

The s framework of C-C and C-H bonds is based
on sp2 hybridized C atoms. a. After accounting
for s bonding, an unhybridized p orbital
remains on each C atom, with one electron per p
orbital b. The p orbitals overlap to form p
bonds that form a continuous p electron cloud
above and below the plane of the ring
Delocalized p bonds
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Chemical Bonding Theory

• Summary of Valence Bond Results
• Bonded atoms share one or more pairs of electrons
• At least one s bond exists between each pair of
  bonded electrons.
• s bonds are cylindrically symmetric along the
  internuclear axis and electrons are concentrated
  - localized - between the bonded atoms.
• An appropriate set of hybrid atomic orbitals is
  formed to form s bonds.
• The set of hybrid orbitals depends on the number
  of s bonds to be formed, the number of
  nonbonded electron pairs and the geometry of the
  molecule.
• AX2 type sp hybridization linear molecular
  geometry
• AX3 or AX2E type sp2 hybridization trigonal
  planar or bent geometry
• AX4, AX3E or AX2E2 type sp3 hybridization
  tetrahedral, trigonal pyramidal, bent molecular
  geometry.
• AX5, AX4E, AX3E2, AX2E3 type sp3d hybridization
  trigonal bipyramid, see saw, T shape or linear
  molecular geometry.
• AX6, AX5E, AX4E2 type sp3d2 hybridization
  octahedral, square pyramid, square planar
  molecular geometry.

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Chemical Bonding Theory

• Summary of Valence Bond Results
• Atoms sharing more than one pair of electrons
  form p bonds by sideways overlap of p atomic
  orbitals.
• p bonds have a nodal plane containing the
  internuclear axis
• In p bonds, electron density is concentrated
  above and below the nodal plane.
• Molecules with two or more resonance structure
  can have p bonds delocalized over more than two
  atoms.

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Chemical Bonding Theory

• Molecular Orbital Theory (MO theory)
• This method deals with interactions of shared
  electrons in a different way.
• Molecular orbitals are formed in such a way that
  they cover the entire molecule.
• The method used is to form Linear Combinations of
  Atomic Orbitals - LCAOs
• There are 2 such combinations from any two atomic
  orbitals
• yMO a1yAO 1 a2yAO 2
• yMO a1yAO 1 - a2yAO 2
• The as indicate how much of each AO is involved
  in yMO. In our case, the as are 1.
• The 1st principle of MO theory is that the total
  number of molecular orbitals is always equal to
  the total number of atomic orbitals contributed
  by the atoms combined
• Other more complicated LCAOs can be constructed
  as more AOs are combined

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Chemical Bonding Theory

• MO Theory
• Two results obtain from MO theory
• The shapes of the molecular orbitals can be
  determined.
• The energies of the molecular orbitals can be
  determined.
• Example H2 molecule

yMO y1s H1 y1s H2 s1s yMO y1s H1 - y1s
H2 s1s
2nd principle of MO theory the bonding MO is
lower in energy than the antibonding MO
s1s2s1s 0
s1s orbital is a bonding orbital - there is
electron density between atoms s1s orbital is an
antibonding MO - there is a nodal plane
perpendicular to the internuclear axis between
the nuclei
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• Chemical Bonding Theory
• MO Theory
• Energy Level Diagram for H2
• Bond order 1( bonding electrons - antibonding
  electrons)
• Energy Level Diagram for He2 Energy
  Level Diagram for He2
• s1s2 s1s1 B. O. 1
  s1s2 s1s2
  B. O. 0
• paramagnetic

3rd principle of MO theory electrons of a
molecule are assigned to orbitals of
successively higher energy according to the
Pauli exlusion principle and Hunds rule of
maximum spin multiplicity
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Chemical Bonding Theory

• MO Theory
• 2nd Period Homonuclear Diatomic Molecules
• Li2 Be2

s1s2 s1s2 s2s2 B. O. 1 s1s2 s1s2
s2s2 s2s2 B. O. 0
diamagnetic
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• Chemical Bonding Theory
• MO Theory
• 2nd Period Homonuclear Diatomic Molecules
• LCAOs for 2p orbitals
• Head-to-head LCAOs
• Sideways overlap

There is a 2nd set of p orbitals perpendicular
to to those shown.
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• Chemical Bonding Theory
• MO Theory
• 2nd Period Homonuclear Diatomic Molecules
• B2 C2

s1s2 s1s2 s2s2s2s2p2p2 B. O. 1 s1s2
s1s2s2s2s2s2p2p4 B. O. 0
paramagnetic
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• Chemical Bonding Theory
• MO Theory
• 2nd Period Homonuclear Diatomic Molecules
• N2 O2

s1s2 s1s2 s2s2s2s2p2p4s2p2 B. O. 3 s1s2
s1s2 s2s2s2s2s2p2p2p4p2p2 B. O. 2
diamagnetic paramagnetic
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• Chemical Bonding Theory
• MO Theory
• 2nd Period Homonuclear Diatomic Molecules
• F2 Ne2

s1s2 s1s2 s2s2s2s2s2p2p2p4p2p4 B. O. 1 s1s2
s1s2 s2s2s2s2s2p2p2p4p2p4s2p2 B. O.0
diamagnetic
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Chemical Bonding Theory

• MO Theory
• Heteronuclear diatomics
• The MO treatment is very similar to that for the
  homonuclear diatomics
• Use the same energy level diagram for F2, except
  that the more electronegative element has AOs
  at lower energy than the less electronegative
  element
• Example CO

• s1s2 s1s2 s2s2s2s2s2p2p2p4
• B. O. 3
• Diamagnetic

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Chemical Bonding Theory

• MO Theory
• Delocalization of p electrons - the MO analogy to
  resonance
• In species such as O3, NO2-, CO32-, benzene,
  resonance was used in valence bond theory to
  explain the equivalence of the multiply bonded
  atoms
• MO theory makes use of electron delocalization of
  the p electrons to explain the same observation

• Example O3
• Assume the central atom is trigonal planar
• The central atom makes 2 s bonds with the
  terminal Os
• Each atom has 3 unused p orbitals having p
  symmetry containing 4 electrons
• The 3 p orbitals form 3 MOs a p bonding MO, a p
  antibonding MO and a p nonbonding MO
• The 4 electrons form p MO electron configuration
  p2 pnb2 pp0
• The p bond order is 0.5, giving a total average
  O-O bond order of 1.5

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Chemical Bonding Theory

• MO Theory
• Delocalization of p electrons - the MO analogy to
  resonance
• Example Benzene

• The six p orbitals on C having p symmetry form
  6 p MOs whose energy diagram is shown
• The three p bonding MOs extend over the 6
  carbons in the ring resulting in electron
  delocalization over the ring of carbon