Chemical Bonding

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Chemical Bonding

• Mr. Doddridge

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Formal Charge

• The formal charge of any atom in a molecule is
  the charge the atom would have if all the atoms
  in the molecule had the same

electronegativity.
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Formal Charge

• Rules

1. All unshared (nonbonding) electrons are assigned
   to the atom on which they are found.
2. For any bond single, double, or triple half
   of the bonding electrons are assigned to each
   atom in the bond.
3. The Formal Charge of each atom is calculated by
   subtracting the number of electrons assigned to
   the atom from the number of valence electrons in
   the isolated atom.

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Formal Charge
2
½(2)
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Formal Charge
2
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Formal Charge
2

• Total electrons assigned to the first oxygen atom
  is 7.
• How about the other two atoms?
• 4 and 5 respectively.
• How do these charges compare to the expected
  charges?
• Since carbon is usually four, its formal charge
  in this configuration is 0 (net change).
• What are the formal charges of the two oxygens?
• Normal charge for oxygen is 6 therefore the
  formal charge of the first oxygen is -1 and the
  second oxygen is?
• 1. (Note the net charge of the molecule.)
• But is this the only possible arrangements of the
  electrons?
• What about?

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Formal Charge


O
O
C

• What are the formal charges on each atom?
• 0, 0, 0
• 22½(4)6 6 0

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Resonance Structures


O
O
C

• Since both structures are possible they are
  called resonance structures.
• In Fact

• Which structure is most likely?

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Molecular GeometriesHybridized OrbitalsMultiple
Bonds
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These can hybridize to form
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• The truth is the one s and two p orbitals
  have hybridized into sp orbitals in fact sp2.

• What shape can this molecule have? (Remember,
  this is simply the central atom).

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• How many p-orbitals are there?
• What is the maximum sp hybridization?
• What would it look like?

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• Notice the angle between all three hybridized
  orbitals will be the same?
• 109.5

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Some Definitions

• Bonds formed from the overlap of orbitals in the
  plane between two atoms are called sigma (s)
  bonds.
• Bonds formed from the overlap of unhybridized p
  orbitals (above and below the central plane) are
  called pi (p) bonds.

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• In this atom we have three sp2 orbitals and one
  unhybridized p orbital.
• The unhybridized p orbital is available to make a
  p bond outside the center plane of the s bond.

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Sigma bond

• On the center plane

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Pi bond

• Starts out looking like this and then the
  unhybridized p orbitals overlap.

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Pi bond
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Pi bond

• This is one pi bond in the y plane.
• This results in one double bond.

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Molecular Geometries
or
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H2O
O
H
H
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H2O
and
or
O
H
H
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The End