AP CHEMISTRY
TOPIC 5: BONDING, PART B
Bond Energy
EXAMPLES
Day 50:
Formal Charge
Bond Energy deals with the energy to make (form) or break a bond between atoms in a molecule. Also, Bond Energy
dictates the strength of a bond – higher bond energy = stronger bonds.
The breaking of a bond requires energy – endothermic ( Positive Energy )
When a bond is formed, energy is released – exothermic ( Negative Energy )
Formal Charge (of an atom) in a molecule is the difference between the number of valence electrons on the free atom and
the number of electrons assigned to the atom in the molecule.
Therefore, to determine formal charge of a given atom in a molecule, you need to know two things:
1) The number of valence electrons on the free neutral atom (which has zero net charge because the number of
electrons equals the number of protons.
2) The number of valence electrons “belonging” to the atom in the molecule.
Formal Charge = ( # of valence e – on free atom) – ( # of electrons assigned to the atom)
“Assigned electrons” may be confusing if we do not understand what is meant by assigned.
( Electrons )ASSIGNED = ( # of lone (un-shared) electrons ) + [ ½ ( number of shared electrons)]
Example: SO4-2 has a total of 32 electrons. In the past we would simply place
electrons around the structure and be happy with what we had done.
Now with formal charges, we want to evaluate if this is the BEST
Lewis model that we can draw, and Formal Charges allow us to do
this evaluation. An element (within the Lewis structure) should try
to achieve a formal charge AS CLOSE TO ZERO ( 0 ) as possible
when drawn correctly. Lets check this drawing.
Find the formal charge for each structure. Whichever structure has the least amount of formal charge (often the
least number of atoms that have formal charge (zero is the formal charge that all atoms are “trying” to achieve))
is the best structure. ( NOTE: All of the structures should add up to the same charge… zero for a neutral
compound, or the charge of the ion if it's an ion )
Electrons assigned to each (single bonded) oxygen = 6 ( lone electrons ) + ½ x 2 ( shared electrons ) = 7
Formal charge on the (single bonded) oxygen = 6 ( valence e - ) minus 7 ( e – assigned ) = - 1
Electrons assigned to sulfur = 0 ( lone electrons ) +
½
x 8 ( shared electrons ) = 4
Formal charge on sulfur = 6 ( valence e - ) minus 4 ( e – assigned ) = 2
TURN THE PAGE
Now for the second drawing:
Oxygen with Single Bonds:
Electrons assigned to each oxygen
= 6 ( lone electrons ) + ½ x 2 ( shared electrons ) = 7
Formal charge on oxygen = 6 ( valence e - ) minus 7 ( e – assigned ) = - 1
Oxygen with Double Bonds:
Electrons assigned to each oxygen = 4 ( lone electrons ) + ½ x 4 ( shared electrons ) = 6
Formal charge on oxygen = 6 ( valence e - ) minus 6 ( e – assigned ) = 0 (zero)
Sulfur:
Electrons assigned to sulfur = 0 ( lone electrons ) + ½ x 12 ( shared electrons ) = 6
Formal charge on sulfur = 6 ( valence e - ) minus 6 ( e – assigned ) = 0 (zero)
THERFORE, the best Lewis Structure for SO4-2 SHOULD be one of the Resonance structures below:
Let’s do another: Calculate the formal charge on each atom and determine the best Lewis Structure for SO2Cl2