Chapter 8 ppt. 2
Bond Energy, Resonance and
Formal Charges
Breaking Vs. Forming Bonds
 To
break bonds, energy must be added
(required, endothermic, energy is positive)
 To form bonds, energy must be removed (given
off, exothermic, energy is negative)
 Change in enthalpy: DH = sum of the energies
required to break old bonds (positive signs) plus
sum of energies released in formation of new
bonds (negative signs)
Bond Energies
 Energy
stored in bonds can be used to
determine energy accompanying various
chemical reactions.
 Bond dissociation energies = energy required to
break bonds (positive kJ/mol, endothermic)…
reverse for forming bonds. Table 8.4 in book
 DH = all bonds broken – all bonds formed
 Must use a balanced equation!
EXAMPLE
data 8.4 in your book to determine DH
for the following reaction.
CH4(g) + 4F2(g)  CF4(g) + 4HF(g)
 Use
 Answer:
-1932 kJ
Localized Electron Bonding
Model…Highlights
 Defined: A molecule
is composed of atoms that
are bound together by sharing pairs of electrons
using the atomic orbitals of the bonded atoms


Localized to one of the atoms (LONE PAIR)
Localized to the space between atoms (BONDING
PAIR)
LE Model Parts
1. Description of the valence electron
arrangement in the molecule using Lewis
structures
2. Prediction of the geometry of the molecule
using the valence shell electron-pair repulsion
(VSEPR) model
3. Description of the type of atomic orbitals used
by the atoms to share electrons or hold lone
pairs
Exceptions to Octet Rule
 If
valence electron total is odd, the octet rule
doesn’t work
 Some atoms do not require all 8 valence
electrons (or have more than 8)
 These molecules can exist in stable form


Boron: forms compounds where boron has less
than 8 electrons (ex: BF3)
More than 8 only happens with elements in period
3 and beyond ex: SF6
 See
pg. 371 purple box for rules if needed
Resonance
 Molecules
with more than one possible electron
dot structure
 Do not switch back and forth
 Molecules exist as a mixture (hybrid) of the
resonance forms
 Use double headed arrow to signify
Formal Charges
Can’t use oxidation numbers because electrons are
not shared evenly between atoms (electronegativities)
 Atoms can be assigned formal charges using the
following:

Formal Charge = (# valence e- on atom) – (# valence eassigned to the atom in the molecule)
# valence e- assigned = (# lone pair e-) + (½ # shared e-)

Atoms want to have formal charges close to zero
 Negative formal charges should be on the most
electronegative atom
 ***ESTIMATES of charges (not exact charges)
Resonance Example

Draw the Lewis dot
(Localized Electron
Model) structure(s) of
CO32-
Example
 Assign
formal charges to each atom in CO2
 Which is more likely?

Answer: two double bonds b/c all formal charges
are zero
 Draw
all resonance structures and show the
most stable one for SCN
Answer: two double bonds b/c N more
electronegative than S