AP Chemistry: Lecture Notes – Chemical Bonding Chapter 8

Test: Friday, November 5, 2010

  • 8.1 Types of Chemical Bonds
  • 8.2 Electronegativity
  • 8.3 Bond Polarity and Dipole Moments
  • 8.4 Ions: Electron Configurations and Sizes
  • 8.5 Energy in Ionic Compounds
  • 8.6 Partial Ionic Character of Covalent Bonds
  • 8.7 Covalent Bonds
  • 8.9 The Localized Electron Bonding Model
  • 8.10 Lewis Structures
  • 8.11 Exceptions to the Octet Rule
  • 8.12 Resonance

Vocabulary:

  1. Chemical bonding
  2. Covalent bond
  3. Delocalized bonding
  4. Dipole moment
  5. Bond energy
  6. Electronegativity
  7. Ionic bond
  8. Isoelectronic
  9. Nonpolar covalent bond
  10. Polar covalent bond

I. Types of Chemical Bonding

  1. Ionic Bonding-

Example:

  1. Covalent Bonding-

1.

2.

Examples:

  1. Metallic Bonding-

II. Valence Electrons

  1. Definition-
  1. Electron Configurations-
  1. Lewis Dot Structures

III. Ionic Bonding

  1. Predicting Chemical Formulas from Dot Structures
  1. Sodium chloride
  1. Aluminum oxide
  1. Lithium nitride

Ionic Compound / Component Ions / Lattice Energy / Tm (melting point)

B. Lattice Energy-

C.General Properties of Ionic Compounds

1.

2.

3.

4.

  1. Ionic Compounds with Polyatomic Ions
  1. Ammonium nitrate
  1. Calcium sulfate
  1. Sodium phosphate
  1. Sizes of Ions
  1. Cations-
  1. Anions-
  1. Isoelectronic Ions -

III. Covalent Bonding

  1. Octet Rule –
  1. Exception #1: (duet rule)
  1. Fluorine gas
  1. Fluorine with Chlorine
  1. Hydrogen fluoride
  1. Double and Triple Bonding
  1. Oxygen gas
  1. Nitrogen gas
  1. CN1-
  1. How to Draw a Lewis Dot Structures
  1. Calculate the number of valence electrons in the molecule
  2. Use a pair of electrons to form a bond between each of the atoms
  3. Distribute remaining electrons so each atom has eight (two for H) electrons
  4. Extra electrons go back on the central atom
  1. Examples:
  1. CH4b. H2O
  1. CH2Od. CO2
  1. Exceptions #2 and #3
  1. Exception #2 - Odd number of electrons or less than eight
  2. NO
  3. NO2
  4. BeCl2
  1. Exception #3 - Expanded Octets
  2. PF5
  3. SF6
  1. Resonance
  1. Definition-
  1. Examples – SO2, N2O
  1. Formal Charge
  1. Definition-
  1. Formal charge = (valence electrons of the atom) – (½ of the bonding electrons+ the number of lone pair electrons)
  1. Two fundamental assumptions about formal charges:
  1. A formal charge of ______is best.
  2. A formal charge of either _____ or ______is acceptable.
  1. Explains
  2. why some are more stable with expanded octets
  3. why some molecules have incomplete octets
  4. which resonance structure is more favorable

IV. Polar Covalent and Nonpolar Covalent BONDS

  1. Electronegativity –
  1. Periodic Trends
  2. Down a group
  3. Across the period

  1. Electronegativity Difference
  1. Example: Order the following bonds according to polarity: H—H, O—H, Cl—H, S—S, and F—H.
  1. Dipole Moments
  1. Definition-
  1. Representation of dipolar character
  1. For each of the molecules, show the direction of the bond polarities and indicate which ones have a dipole moment. HCl, Cl2, SO3, CH4, H2S

V. Covalent Bond Energies – calculating the energy change in a reaction

A. Bond Energy –

  1. Tables – pg. 351
  1. Formula:
  1. Examples:

1. H2 (g) + F2 (g)  2 HF (g)

2. 2 C2H2 + 5 O24 CO2 + 2 H2O

3. CH4 + 2 Cl2 + 2 F2  CF2Cl2

E. Bond Length

F. Bond Order