Bonding:

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Bonding General Concepts Continued
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Covalent Bonding Model

• Remember chemical bonds can be viewed as forces
  that cause a group of atoms to behave as a unit.
• Bonds result from the tendency of a system to
  seek its lowest possible energy.
• Bonds occur when collections of atoms are more
  stable (lower in energy) than the separate atoms.

3
Covalent Bond energies and Chemical Reactions

• Bond Energies
• Bond energy the energy required to break a
  given chemical bond.
• To break bonds, energy must be added to the
  system (endothermic).
• To form bonds, energy must be released
  (exothermic).

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Single bond one pair of electrons
shared. Double bond two pairs of electrons
shared. Triple bond three pairs of electrons
shared.
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• Shared Electron Pairs and Bond Length
• As the number of shared electrons increases, the
  bond length shortens.

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Bond Energy and Enthalpy

• Bond energy values can be used to calculate
  approximate energies for reactions.
• Example calculate the change in energy that
  accompanies the following reaction
• H2 (g) F2 (g) ? 2HF (g)
• To form HF, one H-H bond and one F-F bond must be
  broken and two H-F bonds must be formed.

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• Remember for bonds to be broken energy must be
  added to the system an endothermic process
  and carries a positive sign.
• Formation of a bond releases energy an
  exothermic process and carries a negative sign.
• Enthalpy change
• ?H ?nD(bonds broken) ?nD(bonds formed)
• where ? represents the sum of terms and D
  represents the bond energy per mole (n) of bonds.
• D always has a positive sign.

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In the case of the formation of HF, ?H DH-H
DF-F 2DH-F (1 mol x 432 kJ/mol) (1
mol x 154 kJ/mol) -
(2 mol x 565 kJ/mol) -544 kJ Thus, when
1 mol H2 (g) and 1 mol F2 (g) react to form
2 mol HF (g), 544 kJ of energy should be
released. When this result is compared to the
result for the reaction when using the standard
enthalpy of formation for HF (-542 kJ) the use of
bond energies works well.
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The Covalent Chemical Bond A Model
Model

• Models are attempts to explain how nature
  operates on the microscopic level based on
  experiences in the macroscopic world.

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Fundamental Properties of Models

1. A model does not equal reality.
2. Models are oversimplifications, and are therefore
   often wrong.
3. Models become more complicated and are modified
   as they age.
4. We must understand the underlying assumptions in
   a model so that we dont misuse it.
5. When a model is wrong, we often learn much more
   than when it is right.

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Localized Electron Bonding Model

• A molecule is composed of atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms.

• Electron pairs are assumed to be localized on a
  particular atom or in the space between two
  atoms
• Lone pairs pairs of electrons localized on an
  atom
• Bonding pairs pairs of electrons found in the
  space between the atoms

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Localized Electron Bonding Model has three parts

1. Description of valence electron arrangement
   (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to share
   electrons or hold lone pairs.

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Lewis Structure
G. N. Lewis (1875-1946)

• Shows how valence electrons are arranged among
  atoms in a molecule.
• Reflects central idea that stability of a
  compound relates to noble gas electron
  configuration.

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Duet Rule

• Hydrogen forms stable molecules where it shares
  two electrons.

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Octet Rule

• Elements form stable molecules when surrounded by
  eight electrons.

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Steps for Writing Lewis Structures

1. Sum the valence electrons from all the atoms.
2. Use a pair of electrons to form a bond between
   each pair of bound atoms.
3. Atoms usually have noble gas configurations.
   Arrange the remaining electrons to satisfy the
   octet rule (or duet rule for hydrogen).

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Steps for Writing Lewis Structures

• Sum the valence electrons from all the atoms.
  (Use the periodic table.)
• Example H2O
• 2 (1 e) 6 e 8 e total

• Use a pair of electrons to form a bond between
  each pair of bound atoms.
• Example H2O

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• Atoms usually have noble gas configurations.
  Arrange the remaining electrons to satisfy the
  octet rule (or duet rule for hydrogen).
• Examples H2O, PBr3, and HCN