Chemical Bonding Theory

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• Chemical Bonding Theory
• Valence bond theory is one of two methods of
  viewing how electrons are shared in covalent
  bonding.
• The quantum mechanical approach to valence bond
  theory is that the wave function associated with
  the shared electrons is made up from the atomic
  orbitals on the two bonded atoms so that their
  identity is retained.
• Electrons are localized in the region where the
  bond forms
• The atomic orbitals overlap so as to give a
  maximum in their overlap and put as much
  electron density as possible between the bonded
  atoms.
• This is consistent with the Lewis model which
  places the bonding electrons between the boned
  atoms
• H2 is a simple example each H atom has a 1s
  electron
• The two electrons are shared equally in each
  atoms 1s orbital

The next slide shows how the potential energy of
the two atoms changes as they are brought closer
together from infinite separation The minimum
potential energy occurs when the nuclei are 74
pm apart
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Chemical Bonding Theory
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• Chemical Bonding Theory
• For the heternuclear diatomic HF, the bond
  results from overlap of the 1s orbital on H and
  the half-filled p orbital on
• HF

In terms of the valence bond theory, the bond is
formed by pairing the 1s electon from H with the
2p electron from F to form the electron pair bond.
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• Chemical Bonding Theory
• Valence bond theory
• For H2O, one valence bond picture is that the 1s
  electrons on each H atom overlaps with two
  half-filled p orbitals on O to form two electron
  pair bonds.

For clarity, only the two 2p orbitals on O
involved in bonding are shown. There is also a 2s
and a third 2p valence orbital on O, each with a
pair of electrons. Note this picture predicts a
90o H-O-H bond angle in water. The actual bond
angle is 104.5o and the deviation could come from
the d charges on H due to the electronegativity
difference between H and O.

• For NH3 a similar picture gives three electron
  pair bonds from overlap of the three 1s
  electrons on each H atom and the three 2p
  orbitals on N each with one unpaired electron.

• N 1s22s22px12py12pz1
• This picture predicts
• The H-N-H bond angle of 90o giving a trigonal
  pramidal structure.
• The non-bonded valence electron pair is in a 2s
  orbital

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• Chemical Bonding Theory
• Valence bond theory
• The bonding in carbon because the valence shell
  electron configuration is 2s22px12py1, we would
  expect the simplest compound between C and H
  would be CH2.
• This compound is known but is extremely reactive.
• CH4 is the compound between H and C with one atom
  of C per molecule.
• One way to explain this is to postulate that an
  excited electronic state of carbon forms

• Hybrid orbitals are formed by mixing the 2s
  orbital with the three 2p orbitals. These four
  new orbitals are degenerate

• The hybrid sp3 orbitals can form four CH bonds.
  These bonds point to the corners of a regular
  tetrahedron.

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• Chemical Bonding Theory
• Valence bond theory
• The bonding in carbon
• One advantage of this scheme is that four bonds
  are formed between C and H instead of two bonds.
  Bond formation is exothermic and produces a more
  stable state for carbon and hydrogen. This
  process more than compensates for the energy
  required to form the hybrid orbitals.
• The four new sp3 orbitals are one fourth s and
  three-fourths p in character and are fatter
  than a p orbital.
• Each sp3 orbital has a nodal plane containing the
  nucleus. The lobes are not symmetrical in
  size like a p orbital.

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sp3 hybrid orbitals a. A single sp3 hybrid
orbital showing the two regions of electron
density. b. The four sp3 hybrid orbitals are
directed at the corners of a tetrahedron.
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• Chemical Bonding Theory
• Valence bond theory
• H2O revisited if the O is hybridized sp3,
• There are 2 electron pairs in two sp3 orbitals
  and two unpaired electrons in the other two sp3
  orbitals
• This allows for the formation of two bonds
  between H and O
• The H-O-H bond angle is predicted to be 109.5o,
  but its found to be 104.5o
• The decrease of 5.0o is due to non-bonded
  electron pair - bonded electron pair repulsions
  from the two pair of non-bonded electrons.
• This is easier to explain than the 14.5o increase
  from the earlier model not involving orbital
  hybridization.

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• Chemical Bonding Theory
• Valence bond theory
• NH3 revisited if the N is hybridized sp3,,
• There is one electron pair in one of the sp3
  orbitals and three unpaired electrons in the
  other three sp3 orbitals.
• This allows for the formation of three bond
  between H and N
• The H-N-H bond angle is predicted to be 109.5o,
  but its found to be 107o
• The decrease is only 2.5o caused by repulsion
  between the non-bonded electron pair and the
  bonding pairs of electrons.
• NH3 is a good base indicating the non-bonded
  electron pair is available for donation to
  acids. This would be difficult if this pair were
  in an s orbital on N.

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• Chemical Bonding Theory
• Valence bond theory
• BF3 Only three electron pair bonds are formed.
• The orbital hybridization scheme produces three
  electrons in three equivalent sp2 orbitals.
• Overlap between each sp2 orbital and a p orbital
  in F with one unpaired electron produces three
  electron pair bonds.
• The three sp2 orbitals point to the corners of a
  planar triangle.

For clarity, the non-bonding electrons on F are
not shown Note, there is a left over,
unhybridized p orbital on B. When BF3 reacts
with NH3, the NH3 provides the electrons for a
coordinate- covalent bond. In this case, B will
rehybidize to four sp3 orbitals
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• Chemical Bonding Theory
• Valence bond theory
• BeCl2 Only two electron pair bonds are formed.
• The orbital hybridization scheme produces three
  electrons in two equivalent sp orbitals.
• Overlap between each sp orbital and a p orbital
  in Cl with one unpaired electron produces three
  electron pair bonds.
• The three sp orbitals point in a strait line.

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• Chemical Bonding Theory
• Valence bond theory
• PCl5 sp3d

• SF6 sp3d2

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• Chemical Bonding Theory
• Single bonds in the valence bond theory, single
  bonds are made up of atomic orbitals that are
  cylindrically symmetric about the line joining
  the bonded atoms.
• Such bonds are called sigma - s - bonds.
• Overlap of 2 s orbitals in H2
• Overlap of an s and a p orbital in HF
• Overlap of 2 p orbitals in F2
• Overlap of an s or p orbital with an spy hybrid
  orbital - BeCl2, CH4, PCl5, etc.
• Multiple bonds the second bond involves overlap
  of two p orbitals on different atoms that are
  perpendicular to the internuclear axis

The internuclear axis contains a nodal
plane Electron density is above and below the
nodal plane These bonds are called pi - p - bonds
Single bonds in valence bond theory are s
bonds. Double bonds in valence bond theory are
one s bond and one p bond. Triple bonds in
valence bond theory are one s bond and two p
bonds.
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• Chemical Bonding Theory
• Multiple bond examples
• Ethylene

The 2 s bond from each C to H involve overlap of
C sp2 and H s orbitals The single s bond between
each C involves overlap of C sp2 orbitals The
single p bond between each C involve p overlap of
C unhybridized p orbitals
2 lobes of p bond
The p bond locks this molecule into a planar
structure all 6 atoms are in the same plane.
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• Chemical Bonding Theory
• Multiple bonds example acetylene

The s bonds between C and H involve overlap of C
sp and H s orbitals The s bond between C and C
involve overlap of C sp orbitals The p bonds
between each C involve overlap of two pairs of
unhybridized p orbitals.
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