Chang Chapter 10 Outline

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Chang Chapter 10 Outline

1. Molecular Geometry
2. Dipole Moments
3. Valence Bond Theory
4. Hybridization of Atomic Orbitals
5. Hybridization in Molecules Containing Double and
   Triple Bonds
6. Molecular Orbital Theory

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10.1 Molecular Geometry

• Valence Shell outermost electron-occupied shell
  of an atom, it hold the electrons that are
  usually involved in bonding.
• Valence Shell Electron Pair Repulsion Model
  (VESPR)
• As far as VESPR is concerned, double and triple
  bonds can be considered as single bonds with
  regard to shape
• If a molecule has 2 or more resonance structures,
  we can apply the VESPR model to any one of them

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Ideal Geometries

• There is a fundamental geometry that corresponds
  to the total number of electron pairs around the
  central atom 2, 3, 4, 5 and 6

linear
trigonal planar
tetrahedral
trigonal bipyramidal
octahedral
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Molecules in which the Central Atom has no lone
pairs

• Linear (BeF2) Trigonal Planar (BF3)
• 180 120 (Equilateral Triangle)

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Four Electron Pairs

• Tetrahedral (CH4)
• Bond angles are 109.5

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Five Electron Pairs

• Trigonal bipyramid (PCl5)
• Bond angles vary
• In the trigonal plane, 120
• Between the plane and apexes, 90
• Between the central atom and both apexes, 180

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Six Electron Pairs

• Octahedron (SF6)
• Octahedral or square bipyramid
• Bond angles vary
• 90 in and out of plane
• 180 between diametrically opposite atoms and the
  central atom

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Molecular Geometry Summarized
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Molecular Geometry Summarized
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Molecular Geometry Summary
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Molecular Geometry Summary
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Lone Pairs and Expanded Octets

• Where expanded octets are possible, place the
  extra lone pairs on the central atom
• Example XeF4

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Bond Angles and Lone Pairs

• Ammonia and water show smaller bond angles than
  predicted from the ideal geometry
• The lone pair is larger in volume than a bond
  pair
• There is a nucleus at only one end of the bond so
  the electrons are free to spread out over a
  larger area of space

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One and Two Lone Pairs - Ammonia and Water
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VESPR Guidelines

1. Write the Lewis structure of the molecule
   considering only the electron pairs around the
   central atom
2. Count the number of electron pairs around the
   central atom. Treat double and triple bonds as
   though they were a single bond
3. Predict the overall arrangement of the electron
   pairs
4. Predict the overall geometry of the molecule
5. When predicting bond angles, remember that lone
   pairs repel one another more than a bonding pair
   repels another bonding pair.

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10.2 Dipole Moments
Where µ is the product of the charge q and the
distance r between the charges
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Diatomic Molecules

• Diatomic molecules of the same element are purely
  covalent, have no polarity, and are called
  nonpolar molecules
• Diatomic molecules of different elements have
  dipole moments and are called polar molecules

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The Hydrogen Molecule
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Covalent Bonding Introduction

• Electron density
• Electrons are located between the nuclei of atoms
• When two hydrogen atoms come together, electron
  density is spread over the entire molecule

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Polarity of Molecules
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10.3 Valence Bond Theory

• Valence bond theory (VB) assumes that the
  electrons in a molecule occupy atomic orbitals of
  the individual atoms atoms retain their own
  electrons
• Molecular orbital theory (MO) assumes the
  formation of molecular orbitals from the atomic
  orbitals
• Neither theory perfectly explains all aspects of
  bonding, but each has contributed something to
  our understanding of many observed molecular
  properties
• Atoms attract one another until the repel one
  another (nuclei and electrons included)
• The formation of bonds gives off heat (ouch)

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10.4 Hybridization of Atomic Orbitals

• sp3 hybridization
• C has 2s2 and 2p2 valence electrons 4 total
• 1 paired s electron jumps to the p orbital so we
  now have 1 unpaired s and 3 unpaired p orbitals
• Valence Bond Theory says we now have a
  hypothetical sp3 hybridization (1 from an s and 3
  from a p, making 4 sp3 hybridized orbitals

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sp2 hybridization

• sp2 hybridization
• BF3
• B has 2 paired 2s electrons and 1 unpaired 2p
  electron
• If 1 of the paired 2s electrons is promoted to
  the 2p orbital, we would now have 1 unpaired s
  electron and 2 unpaired p orbitals or a
  hybridized sp2 orbital

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Procedure for Hybridizing Atomic Orbitals

1. Draw the Lewis Dot Structure
2. Predict the overall arrangement of the electron
   pairs (both bonding and lone pairs) using the
   VESPR model
3. Deduce the hybridization of the central atom by
   matching the arrangement of the electron pairs
   with those of the hybrid orbitals shown

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Hybridized Orbitals

• Orbitals Hybrid New Orbitals Degrees
  Compound
• s,p sp 2 orbitals 180 BeCl2
• s,p,p sp2 3 orbitals 120 BF3
• s,p,p,p sp3 4 orbitals 109.5 CH4
• s,p,p,p,d sp3d 5 orbitals 90 and 120 PCl5
• s,p,p,p,d,d sp3d2 6 orbitals 90 and 90 SF6

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10.5 Hybridization of Molecules Containing Double
and Triple Bonds

• Sigma bonds (s) are covalent bonds formed by 2
  orbitals overlapping end to end with the electron
  density concentrated between the nuclei of the
  bonding atoms
• Pi bonds (p) are covalent bonds formed by
  sideways overlapping p orbitals with electron
  densities concentrated above and below the plane
  of the nuclei of the bonding atoms

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10.6 Molecular Orbital Theory

• Molecular orbitals result from interaction of the
  atomic orbitals of the bonding atoms and are
  associated with the entire molecule
• A bonding molecular orbital has lower energy and
  greater stability than the atomic orbitals from
  which it was formed
• An antibonding molecular orbital has higher
  energy and lower stability than the atomic
  orbitals from which it was formed.
• In a sigma molecular orbital, the electron
  density is concentrated symmetrically around a
  line between the two nuclei of the bonding atoms

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Pi Bonds

• In a pi molecular orbital, the electron density
  is concentrated above and below a line joining
  the two nuclei of the bonding atoms.

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Rules Governing Molecular Electron Configuration
and Stability

1. The number of molecular orbitals formed is always
   equal to the number of atomic orbitals combined
2. The more stable the bonding molecular orbital,
   the less stable the corresponding antibonding
   molecular orbital
3. The filling of molecular orbitals proceeds from
   low to high energies.
4. Like an atomic orbital, each molecular orbital
   can accommodate up to two electrons with opposite
   spins
5. When electrons are added to molecular orbitals of
   the same energy, the most stable arrangement is
   predicted by Hunds Rule
6. The number of electrons in the molecular orbitals
   is equal to the sum of all the electrons on the
   bonding atoms.

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Bond Order