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The highlights are:
1. Resonance
2. Possible structures of nitrate
3. Why Resonance Model?
4. Odd electron molecules
5. Formal Charge
1.
Resonance & Formal
Charge
Chapter 8 part 4
2.
What if more than one
valid Lewis dot
structure is possible?
Consider Nitrate ion.
Nitrogen bound to 3
oxygen atoms, one with
a double bond.
But is that the only
possible structure that
obeys the octet rule?
3.
Possible structures for NO3 -
4.
Another illustration
The readers of the
AP do not recognize
this method of
writing the
resonance structure,
but I think it is
important to be
familiar with it.
6.
Why the Resonance model ?
A word about bond
length:
Experimentally it has
been shown:
Single bonds are longer
than double bonds
Double bonds are
longer than triple
bonds.
But…
7.
The Resonance model explains
this…
Since there is more than
one correct position for
the double bond, the
model for resonance
allows that double bond
to shift locations. The
resonance bond is shorter
than a single bond and
larger than a double
bond.
8.
Which of the two
resonance
structures has the
shorter bond
length and …
Why?
9.
Odd Electron Molecules
Relatively few
molecules formed
with nonmetals
contain odd numbers
of electrons.
These examples
usually involve N.
These are not very
stable.
10.
Formal Charge
The difference between the number of
valence electrons on the free atom and the
number of valence electrons assigned to
the atom in the molecule.
Need to determine the valence electron of
the free atom.
Need to determine the number of valence
electrons assigned to the atom in the
molecule.
12.
Formal charge
Assigning electrons:
Lone pairs and single electrons on the
atom each count for that atom.
Bonded electrons are shared and therefore
half go to each atom bound.
Subtract this number form the valence
number.
The result and its sign is the atom’s formal
charge.
13.
In The Sulfate ion,
what are the possible
Lewis dot structures?
Which one is the most
likely?
Use formal charge.
14.
Final example:
Give the possible Lewis dot structures for
XeO3. Which one is the most likely given
the formal charge?
15.
Summary
To calculate the formal charge of an atom:
1. Take the sum of the lone pair electrons
and one half the shared electrons of the
atom in the molecule.
2. Subtract the number of valence electrons
on the free neutral atom.
16.
Summary continued
The sum of the formal charge of all atoms
in a given molecule or ion must equal the
overall charge of that species.
If nonequivalent Lewis structures exist for
a species, those with formal charges
closest to zero and with any negative
formal charges on the most
electronegative atoms are considered to
best describe the bonding in the molecule.
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