Molecular Geometry and
Bonding Theories
Chapter 9 Part 3
November 18th, 2004
32
Hybrid Orbitals
• Consider the CH4 molecule, which has a tetrahedral shape.
• C has the electron configuration 1s22s22p2 and H has 1s1.
Carbon has only 2 unpaired electrons.
• The 2p orbitals in C are at right angles to each other, the 1s
orbital in H is spherical.
• According to Pauling’s theory of orbital hybridization, a new
set of orbitals called hybrid orbitals can be created by
combining the appropriate # of s, p, d and f orbitals.
• This combination results in hybrid orbitals with a spatial
orientation that matches the geometry of the molecule.
• The number of hybrid orbitals is the same as the number of
atomic orbitals used in their construction.
construction.
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Hybridization in Methane
• By mixing the (one) 2s and (three) 2p valence
orbitals in the C atom a new set of hybrid
orbitals called the (four) sp3 orbitals is created.
• The four sp3 hybrid orbitals are tetrahedral in
shape.
• All four sp3 orbitals have identical shapes and
are equivalent in energy.
• Each H atom bonds to the C atom by orbital
overlap between the 1s orbital and the sp3
hybrid orbitals.
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1
sp2 Hybrid Orbitals
• sp2 hybrid orbitals are formed by the mixing of
one s and two p orbitals.
• The large lobes of sp2 hybrid orbitals lie in a
trigonal plane and the angles between them
are 120°.
• All molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central
atom.
• One unhybridized p orbital is left as a result of
the formation of sp2 orbitals.
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36
sp Hybrid Orbitals
• Formed by the combination of one s and one
p orbital.
• The lobes of sp hybrid orbitals are 180º apart,
meaning the geometry is linear.
• All molecules with linear geometry will have sp
hybrid orbitals on the central atom.
• The formation of sp hybrid orbitals leaves 2
unhybridized p orbitals on the central atom.
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2
38
sp3d and sp3d2 Hybrid Orbitals
• Since there are only three p-orbitals, trigonal
bipyramidal and octahedral electron pair
geometries must involve d-orbitals.
• Trigonal bipyramidal electron pair geometries
require sp3d hybridization.
• Octahedral electron pair geometries require
sp3d2 hybridization.
• The electron pair geometry from VSEPR
theory always determines the hybridization.
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40
3
41
Deciding Hybridization
• Decide the hybridization of the central atom
for SF3+ and I3- ions.
I
+
F
S
I
F
I
F
4 electron pairs around the
central atom = tetrahedral.
5 electron pairs around the
central I atom = trigonal
bipyramidal
Hybridization of the S atom =
sp3
Hybridization of the I atom =
sp3d
42
Multiple Bonds
• σ-bonds: electron density lies on
the axis between the nuclei. All
single bonds are σ-bonds.
• π-bonds: electron density lies
above and below the plane of the
nuclei.
• A double bond consists of one σbond and one π-bond.
• A triple bond has one σ-bond and
two π-bonds.
• Often, the p-orbitals involved in πbonding come from unhybridized
orbitals.
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4
Multiple Bonds
• Ethylene, C2H4, has:
– one σ- and one π-bond
– both C atoms are sp2 hybridized
– both C atoms have trigonal planar electron pair and
molecular geometries
– The formation of a π bond requires that the molecule be
planar.
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45
Multiple Bonds
Consider acetylene, C2H2
–The electron pair geometry of each C is linear & the C atoms
are sp hybridized;
–The sp hybrid orbitals form the C-C and C-H σ-bonds;
–There are two unhybridized p-orbitals which form the two πbonds.
–one π-bond is above and below the plane of the nuclei;
–one π-bond is in front and behind the plane of the nuclei.
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5
Delocalized π Bonding
• So far all the bonds we have encountered are
localized between two nuclei.
• In the case of benzene
– there are 6 C-C σ bonds, 6 C-H σ bonds,
– each C atom is sp2 hybridized,
– and there are 6 unhybridized p orbitals on each C atom.
47
Delocalized π Bonding
• In benzene there are two options for the 3 π bonds
– localized between C atoms or
– delocalized over the entire ring (i.e. the π
electrons are shared by all 6 C atoms).
• Experimentally, all C-C bonds are the same length in
benzene.
• Therefore, all C-C bonds are of the same type (recall
single bonds are longer than double bonds).
48
Multiple Bonds
• Every two atoms share at least 2 electrons.
• Two electrons between atoms on the same
axis as the nuclei are σ bonds.
• σ-bonds are always localized.
• If two atoms share more than one pair of
electrons, the second and third pair form πbonds.
• When resonance structures are possible,
delocalization is also possible.
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6
The MO Theory
• Some aspects of bonding are not explained by Lewis
structures, VSEPR theory and hybridization. (E.g.
why does O2 interact with a magnetic field?; Why are
some molecules colored?)
• For these molecules, we use Molecular Orbital
(MO) Theory.
• Just as electrons in atoms are found in atomic
orbitals, electrons in molecules are found in
molecular orbitals.
50
Molecular Orbitals
•
•
•
•
Each MO contains a maximum of two electrons;
MOs have definite energies;
MOs can be visualized with contour diagrams;
MOs are associated with an entire molecule.
51
The Hydrogen Molecule
• When 2 H atoms bond, 1s (H) + 1s (H) must
result in two MOs for H2:
– one has electron density between nuclei (bonding
MO);
– one has little electron density between nuclei
(antibonding MO).
• MOs resulting from s orbitals are σ MOs.
• Bonding (σ) MOs have lower energy than
antibonding (σ*) MOs.
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7
The Hydrogen Molecule
53
The MO Diagram
• Energy level diagram or MO diagram shows
the energies and electrons in an orbital.
• The total number of electrons in all atoms are
placed in the MOs starting from lowest energy
(σ1s) and ending when you run out of
electrons.
• Note that electrons in MOs have opposite spins.
• H2 has two bonding electrons.
• He2 has two bonding electrons and two
antibonding electrons.
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MO Diagram
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8