Molecular Geometry and Bonding Theories

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Molecular Geometry and Bonding Theories
David P. White University of North Carolina,
Wilmington
Chapter 9
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Molecular Shapes

• Lewis structures give atomic connectivity they
  tell us which atoms are physically connected to
  which.
• The shape of a molecule is determined by its bond
  angles.
• Consider CCl4 experimentally we find all Cl-C-Cl
  bond angles are 109.5?.
• Therefore, the molecule cannot be planar.
• All Cl atoms are located at the vertices of a
  tetrahedron with the C at its center.

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Molecular Shapes
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Molecular Shapes
In order to predict molecular shape, we assume
the valence electrons repel each other.
Therefore, the molecule adopts whichever 3D
geometry minimized this repulsion. We call this
process Valence Shell Electron Pair Repulsion
(VSEPR) theory.
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The VSEPR Model
The valence electrons in a molecule are the
bonding pairs of electrons as well as the lone
pairs. There are 11 shapes that are important to
us linear (3 atoms, AB2) bent (3 atoms,
AB2) trigonal planar (4 atoms, AB3) trigonal
pyramidal (4 atoms, AB3) T-shaped (4 atoms,
AB3) tetrahedral (5 atoms, AB4) square planar
(5 atoms, AB4) see-saw (5 atoms, AB4) trigonal
bipyramidal (6 atoms, AB5) square pyramidal (6
atoms, AB5) and octahedral (7 atoms, AB6).
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The VSEPR Model
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The VSEPR Model
Predicting Molecular Geometries To determine the
shape of a molecule, we distinguish between lone
pairs (or non-bonding pairs, those not in a bond)
of electrons and bonding pairs (those found
between two atoms). We define the electron pair
geometry by the positions in 3D space of ALL
electron pairs (bonding or non-bonding). The
electrons adopt an arrangement in space to
minimize e--e- repulsion.
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The VSEPR Model
Predicting Molecular Geometries
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The VSEPR Model
Predicting Molecular Geometries
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The VSEPR Model

• Predicting Molecular Geometries
• To determine the electron pair geometry
• draw the Lewis structure
• count the total number of electron pairs around
  the central atom
• arrange the electron pairs in one of the above
  geometries to minimize e--e- repulsion
• multiple bounds count as one bonding pair

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The VSEPR Model
Predicting Molecular Geometries
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The VSEPR Model
Molecules with Expanded Valence Shells
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The VSEPR Model
Molecules with More than One Central Atom In
acetic acid, CH3COOH, there are three central
atoms. We assign the geometry about each central
atom separately.
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Polarity of Molecules
Polar molecules interact with electric fields. If
the centers of negative and positive charge do
not coincide, then the molecule is polar.
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Polarity of Molecules
If two charges, equal in magnitude and opposite
in sign, are separated by a distance r, then a
dipole is established. The dipole moment, m, is
given by m Qr where Q is the magnitude of
charge. Dipole Moments of Polyatomic
Molecules In a polyatomic molecule, each bond can
be a dipole. The orientation of these individual
dipole moments determines whether the molecule
has an overall dipole moment.
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Polarity of Molecules
Dipole Moments of Polyatomic Molecules Example
in CO2, each C-O dipole is canceled because the
molecule is linear. In H2O, the H-O dipoles do
not cancel because the molecule is bent.
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Polarity of Molecules

• Dipole Moments of Polyatomic Molecules
• It is possible for a molecule with polar bonds to
  be either polar or non-polar.
• For diatomic molecules
• polar bonds always result in an overall dipole
  moment.
• For triatomic molecules
• if the molecular geometry is trigonal pyramidal,
  there is an overall dipole moment
• if the molecular geometry is trigonal planar and
  all three bonds are identical, there is no
  overall dipole moment
• if the molecular geometry is trigonal planar and
  one or two bonds are different, there is an
  overall dipole moment.

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Polarity of Molecules
Dipole Moments of Polyatomic Molecules
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Polarity of Molecules

• Dipole Moments of Polyatomic Molecules
• For tetraatomic molecules with identical bonds
• if the molecular geometry is tetrahedral or
  square planar, then the molecules are nonpolar
• if the molecular geometry is see-saw, the
  molecule is polar.
• For tetraatomic molecules in which one, two, or
  three bonds are different
• the molecule is polar.

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Covalent Bonding and Orbital Overlap

• Lewis structures and VSEPR do not explain why a
  bond forms.
• How do we account for shape in terms of quantum
  mechanics?
• What are the orbitals that are involved in
  bonding?
• We use Valence Bond Theory
• Bonds form when orbitals on atoms overlap.
• There are two electrons of opposite spin in the
  orbital overlap.

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Covalent Bonding and Orbital Overlap
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Covalent Bonding and Orbital Overlap

• As two nuclei approach each other their atomic
  orbitals overlap.
• As the amount of overlap increases, the energy of
  the interaction decreases.
• At some distance the minimum energy is reached.
• The minimum energy corresponds to the bonding
  distance (or bond length).
• As the two atoms get closer, their nuclei begin
  to repel and the energy increases.
• At the bonding distance, the attractive forces
  between nuclei and electrons just balance the
  repulsive forces (nucleus-nucleus,
  electron-electron).

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Hybrid Orbitals

• sp Hybrid Orbitals
• Consider the BeF2 molecule (experimentally known
  to exist)
• Be has a 1s22s2 electron configuration.
• There is no unpaired electron available for
  bonding.
• We conclude that the atomic orbitals are not
  adequate to describe orbitals in molecules.
• We know that the F-Be-F bond angle is 180? (VSEPR
  theory).
• We also know that one electron from Be is shared
  with each one of the unpaired electrons from F.

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Hybrid Orbitals
sp Hybrid Orbitals We assume that the Be orbitals
in the BeF bond are 180? apart. We could promote
an electron from the 2s orbital on Be to the 2p
orbital to get two unpaired electrons for
bonding. BUT the geometry is still not
explained. We can solve the problem by allowing
the 2s and one 2p orbital on Be to mix or form a
hybrid orbital (process called hybridization). Th
e hybrid orbital comes from an s and a p orbital
and is called an sp hybrid orbital.
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Hybrid Orbitals
sp Hybrid Orbitals The two lobes of an sp hybrid
orbital are 180? apart.
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Hybrid Orbitals
sp Hybrid Orbitals Since only one of the Be 2p
orbitals has been used in hybridization, there
are two unhybridized p orbitals remaining on Be.
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Hybrid Orbitals

• sp2 and sp3 Hybrid Orbitals
• Important when we mix n atomic orbitals we must
  get n hybrid orbitals.
• sp2 hybrid orbitals are formed with one s and two
  p orbitals. (Therefore, there is one
  unhybridized p orbital remaining.)
• The large lobes of sp2 hybrids lie in a trigonal
  plane.
• All molecules with trigonal planar electron pair
  geometries have sp2 orbitals on the central atom.

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Hybrid Orbitals
sp2 and sp3 Hybrid Orbitals
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Hybrid Orbitals

• sp2 and sp3 Hybrid Orbitals
• sp3 Hybrid orbitals are formed from one s and
  three p orbitals. Therefore, there are four
  large lobes.
• Each lobe points towards the vertex of a
  tetrahedron.
• The angle between the large lobs is 109.5?
• All molecules with tetrahedral electron pair
  geometries are sp3 hybridized.

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Hybrid Orbitals
sp2 and sp3 Hybrid Orbitals
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Hybrid Orbitals

• Hybridization Involving d 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.
• Note the electron pair geometry from VSEPR theory
  determines the hybridization.

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Hybrid Orbitals
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Hybrid Orbitals
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Hybrid Orbitals

• Summary
• To assign hybridization
• draw a Lewis structure
• assign the electron pair geometry using VSEPR
  theory
• from the electron pair geometry, determine the
  hybridization and
• name the geometry by the positions of the atoms.

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Multiple Bonds
?-Bonds electron density lies on the axis
between the nuclei. All single bonds are
?-bonds. ?-Bonds electron density lies above
and below the plane of the nuclei. A double bond
consists of one ?-bond and one ?-bond. A triple
bond has one ?-bond and two ?-bonds. Often, the
p-orbitals involved in ?-bonding come from
unhybridized orbitals.
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Multiple Bonds
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Multiple Bonds

• Ethylene, C2H4, has
• one ?- and one ?-bond
• both C atoms sp2 hybridized
• both C atoms with trigonal planar electron pair
  and molecular geometries.

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Multiple Bonds

• Consider acetylene, C2H2
• the electron pair geometry of each C is linear
• therefore, the C atoms are sp hybridized
• the sp hybrid orbitals form the C-C and C-H
  ?-bonds
• there are two unhybridized p-orbitals
• both unhybridized p-orbitals form the two
  ?-bonds
• one ?-bond is above and below the plane of the
  nuclei
• one ?-bond is in front and behind the plane of
  the nuclei.
• When triple bonds form (e.g. N2) one ?-bond is
  always above and below and the other is in front
  and behind the plane of the nuclei.

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Multiple Bonds

• Delocalized p Bonding
• So far all the bonds we have encountered are
  localized between two nuclei.
• In the case of benzene
• there are 6 C-C ? bonds, 6 C-H ? bonds,
• each C atom is sp2 hybridized,
• and there are 6 unhybridized p orbitals on each C
  atom.

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Multiple Bonds
Delocalized p Bonding
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Multiple Bonds

• Delocalized p Bonding
• In benzene there are two options for the 3 ?
  bonds
• localized between C atoms or
• delocalized over the entire ring (i.e. the ?
  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).

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Multiple Bonds
Delocalized p Bonding
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Multiple Bonds

• General Conclusions
• Every two atoms share at least 2 electrons.
• Two electrons between atoms on the same axis as
  the nuclei are ? bonds.
• ?-Bonds are always localized.
• If two atoms share more than one pair of
  electrons, the second and third pair form
  ?-bonds.
• When resonance structures are possible,
  delocalization is also possible.