Chapter 6 review and new to Ch 7

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					Chapter 6

  Review
         Types of Chemical Bonding
• Ionic or Covalent?
    – Review: Use electronegativity (p161) differences to
      classify bonding between chlorine and:
        • Calcium
        • Oxygen
        • Bromine

Electroneg diff               Bond type           More–neg atom
   3.0 – 1.0 = 2.0            ionic                   chlorine
   3.5 – 3.0 = 0.5            polar-covalent          oxygen
   3.0 – 2.8 = 0.2            nonpolar-covalent       chlorine

< 0.3 = nonpolar-covalent   between = polar        > 1.7 = ionic
 Characteristics of the Covalent Bond
• Longer bonds tend to be weaker
• Bond strength is relfected by the amount of
  energy needed to break a bond, or the bond
  energy.
The Octet Rule in Covalent Bonding
A. Covalent compounds tend to form so that
   each atom, by sharing electrons, has an octet
   of electrons in its highest occupied energy
   level
  – Exception: H (smallest shell), B (6 electrons),
    (expanded shell): F, O, Cl
Electron-Dot Notation
Lewis Structures
Multiple Covalent Bonds
       **Resonance Structures
• Refers to bonding in molecules or ions that
  can’t be correctly represented by a single
  Lewis structure.
   Formation of Ionic Compounds
• A compound composed of positive and negative
  ions that are combined so that the numbers of
  positive and negative charges are equal
  **Most are crystalline solids
  **Lattice energy is the energy released when one mole
    of an ionic crystalline compound is formed from
    gaseous ions.
  **Lattice energies are listed as negative values to
    indicate that energy is given off when the gaseous
    ions come together to form the ionic crystal
             Polyatomic Ions
• Although polyatomic ions are involved in ionic
  bonding, the ions themselves are held
  together by covalent bonding.
Metallic Bonding

    Section 6.4
      The Metallic-Bond Model
• Metallic bonding results from the attraction
  between metal atoms and the surrounding
  sea of electrons.
  – Electrons are mobile, they can move from the
    region around one nucleus to the region around
    another
     • This allows for the conductivity of heat and electricity
           Metallic Properties
• Shiny
• Malleability – ability of a substance to be
  hammered into sheets
• Ductility – ability of a substance to be drawn
  into wire
        Metallic Bond Strength
• The higher the enthalpy of vaporization, the
  stronger the metallic bonding
  – The amount of energy absorbed as heat when a
    specified amount of a substance vaporizes at
    constant pressure.
Molecular Geometry

     Section 6.5
                VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
• Repulsion between the sets of valence-level
  electrons surrounding an atom causes these sets to
  be oriented as far apart as possible
VSEPR and Unshared Electron Pairs
1. Unshared pairs take up positions in the
   geometry of molecules just as atoms do
2. Unshared pairs have a relatively greater effect
   on geometry than do atoms
3. Lone (unshared) electron pairs require more
   room than bonding pairs (they have greater
   repulsive forces) and tend to compress the
   angles between bonding pairs
4. Lone pairs do not cause distortion when bond
   angles are 120° or greater
              Practice Problems
• Use VSEPR theory to predict the molecular
  geometry of the following molecules:
  a. HI
  b. CBr4
  c. CH2Cl2

  Answers:
  a. Linear
  b. Tetrahedral
  c. Tetrahedral
                       Practice Problems
• Use VSEPR theory to predict the molecular
  geometries of the molecules whose Lewis
  structures are given below
           .. .. ..               .. .. ..
     a.   :F–S–F:            b. : Cl – P – Cl :
          ..   .. ..              ..            ..
                                       : Cl :
                                         ..
Answers
a.    Bent or angular
b.    Trigonal-pyramidal
               Hybridization
• Is the mixture of two or more atomic orbitals
  of similar energies on the same atom to
  produce new hybrid atomic orbitals of equal
  energies.
• Hybrid orbitals are orbitals of equal energy
  produces by the combination of two or more
  orbitals on the same atom.
• The number of orbitals
  involved in hybridization
  determines the
  geometry of a molecule.

• There must be space
  around the atom for
  every hybrid orbital;
  each orbital will be as
  far apart as possible
  from every other
  orbital.
         Intermolecular Forces
• Boiling point is a measure of the force of
  attraction between molecules in covalent
  compounds, between ions in ionic
  compounds, and between atoms in metals.
  Molecular Polarity & Dipole-Dipole
               Forces
• Dipole is created by equal but opposite
  charges that are separated by a short distance
  – The direction of the dipole is from the positive to
    the negative pole
           Hydrogen Bonding
• Intermolecular force in which a hydrogen
  atom that is bonded to a highly
  electronegative atom is attracted to an
  unshared pair of electrons of an
  electronegative atom in a nearby molecule.
      London Dispersion Forces
• The intermolecular attractions resulting from
  the constant motion of electrons and the
  creation of instantaneous dipoles
  – All molecules experience these forces
  – Insignificant relative to dipole-dipole forces
  – They increase with increasing atomic mass

				
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