Formic acid: HCOOH

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• Formic acid HCOOH

Acetone
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• Benzene C6H6

Kekulé structures
Resonance structures each point corresponds to a
CH Each C is sp2 hybridized, one of the sp2
forming a s-bond with H 1s orbital and the other
two forming s-bonds with adjacent C sp2
orbitals.
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• The un-hybridized p orbital on each C is
  available for p-bonding with p orbitals on either
  of the adjacent C atoms

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• Actual structure of benzene is a resonance hybrid
  of the two alternating bond patterns the 6 C
  atoms are identical, and the electrons in the
  p-bonds spread around the entire ring
• This lowers the energy of the molecule -
  resonance adds stability to a molecule

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Characteristics of p bonds bonds Energy of CC
is lt 2 x energy of C-C bond Energy of C?C is lt 3
x energy of C-C bond C, N, O form double bonds
with one another and with elements from later
periods Double bonds are rarely found between
elements in period 3 are below - atoms are too
large for effective side-by-side
overlap. Molecules with alternate double-single
bonds - conjugated molecules
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Isomers Molecules with the same molecular
formula but different structures
cis-1,2-dichloroethylene
trans-1,2-dichloroethylene
Rotation can occur about a single sigma
bond Rotation is restricted about a double bond
isomers are a consequence
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Change of shape triggers a signal along the optic
nerve
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Molecular Orbital Theory

• VB theory localized bond
• VB theory provides the basis of calculating
  electron distributions in molecules but cannot
  explain the properties of some molecules.
• O2
• VB theory
• O Is2 2s2 2p4
• sp2 hybridized O, one sp2 from each forms s-bond
  and the other two are occupied with the lone
  pairs.
• The un-hybridized p on each forms the p-bond
• Indicates that in O2 molecule, all electrons are
  paired.
• However O2 was observed to be paramagnetic

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• VB theory assumes that the electrons are
  localized between the two bonding atoms
• Molecular orbital theory electrons are spread
  throughout the entire molecule electrons are
  delocalized over the whole molecule.
• Pure atomic orbitals combine to produce molecular
  orbitals that are spread out, delocalized, over
  an entire molecule
• Molecular orbitals are built by adding together
  -superimposing - atomic orbitals belonging to the
  valence shell of the atoms in the molecules.

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• H2 wavefunction representing the molecular
  orbitals (MOs) for H2 can be represented by
  combining the two atomic orbitals (AOs) for the
  separated H atoms.
• Wavefunction of the H2 MO
• y yA1s yB1s
• yA1s or yB1s 1s orbital centered on one of the H
  atom(A or B)
• The molecular orbital, y, is a linear combination
  of atomic orbitals
• Any molecular orbital formed from a superposition
  of atomic orbitals is called a LCAO-MO.
• y is a bonding orbital energy of y is lower
  than that of either AO
• In H2, the contribution from each AO to the MO is
  equal

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• The two AOS are waves centered on different
  nucleii.
• Bonding orbital AO wavefunctions interfere
  constructively - MO wavefunction in blue.

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• N AOs overlapping will form N MOs
• Two H AOs overlapping form two Mos one of which
  is the bonding orbital, y.
• The wavefunctions of the two H AOs can also
  interfere destructively - anti-bonding MO of
  higher energy than each of the AOs
• y- yA1s - yB1s
• Node between two nuclei

Probability of finding electrons between nuclei
reduced nuclei repel each other
http//www.shef.ac.uk/chemistry/orbitron/index.htm
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Molecular Orbital Energy Level Diagram
Energy of bonding MO lt AO Energy of
anti-bonding MO gt AO
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• Diatomic Molecules
• Build all possible MOs from available valence AOs
• Then accommodate valence electrons in molecular
  orbitals using the aufbau principles
• 1) Electrons occupy the lowest energy MOs first,
  then orbitals of increasing energy
• 2) Pauli exclusion principle each orbital can
  occupy up to two electrons if two electrons in
  an orbital must be paired
• 3) Hunds rule if more than one orbital of the
  same energy is available electrons enter them
  singly with parallel spinds.

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Lowest unoccupied MO (LUMO)
Highest occupied MO (HOMO)

• H2 molecular orbital energy-level diagram or
  correlation diagram
• Bonding MO - s1s anti-bonding MO s 1s

H2 Ground state electron configuration (s1s)2
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• Bond Order
• 0.5(number of electrons in bonding MOs
• - number of electrons in anti-bonding MOs)

H2 bond order 0.5 (s1s)1
H2 bond order 1 (s1s)2
He2 bond order 0 (s1s)2 (s1s)2
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• Period 2 elements
• In period 2 elements each atom has one 2s and
  three 2p valence AOs expect to form eight MOs
• The two 2s orbitals (one from each atom) overlap
  to form a s2s bonding MO and a s2s antibonding
  MO
• The six 2p orbitals (three from each atom)
  overlap to form six MOs
• The two 2p-orbitals directed toward each other
  form a bonding s-orbital (s2p) and an
  anti-bonding s-orbital (s2p)
• Two 2p orbitals that are perpendicular to the
  internuclear axis overlap side by side to form
  two bonding p and two anti-bonding p orbitals.

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Anti-bonding
Bonding

• s and s orbitals formed from p AOs

p and p orbitals formed from p AOs
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MO diagram for homonuclear diatomic molecules O2
and F2

• MO diagram for homonuclear diatomic molecules Li2
  through N2