Molecular Geometries and Bonding Theories

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Molecular Geometries and Bonding Theories

• Molecular Shapes
• The VSEPR Model
• Polarity of Molecules
• Covalent Bonding and Orbital Overlap
• Hybrid Orbitals
• Multiple Bonds
• Molecular Orbitals
• Second-Row Diatomic Molecules

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Molecular Shapes

• Lewis structures do not indicate the shapes of
  molecules they simply show the number and types
  of bonds between atoms
• The shape of a molecule is determined by it bond
  angles
• Consider CCl4 experimentally we find all Cl-C-Cl
  bond angles are 109.5o
• Thus, the molecule cannot be planar

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• In order to predict molecular shape, we assume
  the valence electrons repel each other.
  Therefore, the molecule adopts whichever 3D
  geometry minimizes 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 as well as the lone pairs
• To determine the shape of a molecule we
  distinguish between lone pairs and bonding pairs
• We define the electron pair geometry by the
  position in 3D space of electron pairs
• The electrons adopt an arrangement in space to
  minimize e--e- repulsion

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• There are only 5 electron pair geometries we
  consider
• Linear
• Trigonal planar
• Tetrahedral
• Trigonal bipyramidal
• Octahedral
• To determine 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 electron repulsion

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• To determine the 3D structure of a molecule
• Draw the Lewis structure
• Place the electron pairs around the central atom
  to minimize repulsion
• Describe the molecular geometry in terms of the
  angular arrangements of the bonding pairs
• A multiple bond is counted as one bonding pair
  when predicting geometry

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• The Effect of Nonbonding Electrons and Multiple
  Bonds on Bond Angles
• Nonbonding electrons exert greater repulsive
  forces on adjacent electron pairs and thus tend
  to compress the angles between the bonding pairs
• Electrons in multiple bonds exert a greater
  repulsive force and adjacent bonds than do single
  bonds

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• Molecules with Expanded Valence Shells
• Atoms which have expanded octets have AB5
  (trigonal bipyramidal) or AB6 (octahedral)
  electron pair geometries
• For trigonal bipyramidal structures, there is a
  plane containing three electron pairs. The fourth
  and fifth electron pairs are located above and
  below the plane
• For octahedral structures, there is a plane
  containing four electron pairs. Similarly, the
  fifth and sixth electron pairs are located above
  and below the plane
• What do you expect for AB7 geometry?

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• Trigonal Bipyramid
• The three electron pairs in the plane are called
  equatorial
• The two electrons above and below the plane are
  called axial
• The axial electron pairs are 180o apart and at
  90o to the equatorial electrons
• The equatorial electrons are 120o apart
• To minimize repulsion, lone pairs are always
  place in equatorial positions

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• Octahedron
• The four electron pairs in the plane are at 90o
  to each other
• The two axial electron pairs are 180o apart and
  90o to the electrons in the plane
• Because of the symmetry of the system, it does
  not matter where the lone pairs are placed

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• Molecules with More Than One Central Atom
• In acetic acid, CH3COOH, there are three central
  atoms two C and one O
• We assign the molecular (and electron pair)
  geometry about each central atom separately
• The first C is tetrahedral
• The second C is trigonal planar, and
• The O is bent (tetrahedral)

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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
• If two charges, equal in magnitude and opposite
  in sign, are separated by a distance d, then a
  dipole is established
• The dipole moment, m, is given by
• m Q.r
• Where Q is the magnitude of the charge

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• 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
• It is possible for a molecule with polar bonds to
  be either polar or nonpolar

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Covalent Bonding and Orbital Overlap

• Lewis structures and VSEPR theory give us the
  shape and location of electrons in a molecule
• They do not explain why a chemical bond forms
• We ask how to account for molecular shape in
  terms of quantum mechanics?. That is, what are
  the orbitals involved in bonding
• We use Valence Bond Theory

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• Valence Bond Theory
• A covalent bond forms when the orbitals on two
  atoms overlap
• The shared region of space between the orbitals
  is called the orbital overlap
• There are two electrons of opposite spin in the
  orbital overlap

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Hybrid Orbitals

• Hybrid orbitals provide a convenient model for
  using valence-bond theory to describe covalent
  bonds in molecules
• To predict the hybrid orbital
• Draw the Lewis structure for the molecule or ion
• Determine the electron-pair geometry using the
  VSEPR model
• Specify the hybrid orbitals needed to accommodate
  the electron pairs based on their geometrical
  arrangements