__
__
?eoi
Strategy for Writing Lewis Structures
1. Total number of valence electrons in the system(Add negative charge iyou
have an anion)
2. Number of electrons if each atom is to be happy.
3. The number of bonds in the system. Sharing electrons in covalent bonds.
4. Draw the structure
Okay! Time for a practice problem!!!!
Example:
.
2
3
Co
1.
=4
C
O=6x3=18
= 2
Charge on ion
24
Total
2. Number of electrons if each atom is to be happy
C needs 8 electrons
0 needs 8 electrons
Electrons for happiness
8
=24
= 32
=
The -2 charge comes as a result of the electrons in the system. The charge is
NEVER counted towards happiness.
3. The number of bonds in the system. Covalent bonds are made by sharing
electrons. You need 32 electrons and you have 24. You are short 8
If you make 4 bonds(with 2 electrons/bond), you will make up the
deficiency. Therefore,
32-24=4 Bonds
2
2
4. Draw the structure. The central atom is carbon. The oxygens go around it!
Because there are 4 bonds, there will be two single bonds and one double
bond. Each bond accounts for 2 electrons. Then complete the octet by putting
electrons around each atom. Double check your results by counting total
electrons in the system..
# bonds
=
#2
-
#1
=
—,
°:
WE ARE HAPPY???? ©
J
Exceptions to the Octet Rule
Some central atoms can exceed the octet rule. This can happen
because there are empty d orbitals that shared electrons can
occupy.
This can happen with third period elements.
For example IC!
3
1. Total of valence electrons
.1=7
• Cl=7(3)=21
• Total = 28 electrons
2. Bonded electrons
Cl—I—Cl
Cl
28-6
=
22 electrons left to fulfill the octet rule
18 electrons can be placed around each Cl atom
Ei—i—
‘‘
:J1:
•‘
There are 4 electrons left. In this scenario you would place any
extra electrons to the central orn and it would look like this:
cl—i—cl
Chapter 8-Section 12- Resonance
Is invoked when more than one valid structure is possible and
can be written for a Lewis Structure. If there are three possible
structures than the resulting electron structure of the molecule
is given by the average of these resonance structures.
Carbonate ion is one example: C0
2
3
Eo
L
I
:Q.
I
L
The LE Model postulates that electrons are localized between a
given pair of atoms. However this does not always operate this
way in nature. Electrons are really delocalized where they can
move around the entire molecule. Resonance is necessary to
compensate for the defective assumption of the LE Model.
FORMAL CHARGES
Since the LE Model operates on pairs of electrons, what if we
have an odd number of electrons?????
Formal charge is the difference between the number of valance
electrons on the free atom and the number of valence electrons
assigned to the atom in a molecule. This is a computed value!
Formal charge is a more realistic estimate of charge
distribution in a molecule.
Process used:
(1) How many electrons each atom “owns” =
(# valence electrons around atom + # bonds)
(2) The formal charge =( # valence electrons on the neutral
atom # electrons owned by the atom based on the
resonance structure)
-
(3) the sum of the formal charges must always equal the
charge on the atom.
Formal Charge and Resonance
Let’s review your textbooks discussion on formal charge and how it is used to draw resonance structures.
Formal charge is, “the difference between the number of valence electrons on the free atom and the number of
valence electrons assigned to the atom in a molecule.” Formal charge is a computational device based on a
localized electron (LE) model and as such is riot perfectly correct. To determine formal charge (a somewhat
more realistic estimate of charge distribution in a molecule), we need to know:
1.
How many electrons each atom “owns.”
[# valence electrons around the atom +
Electrons Owned
4 bonds (which equals Yz 4 shared electrons)]
2.
The formal charge on each atom.
4 valence electrons on the neutral atom
Formal Charge
—4 electrons owned by the atom based on the resonance structure you drew
Lt’s look at CO? again.
2-
;o:b
4 valence electrons on the
•
Carbon owns 4 electrons (4 bonds). The formal charge on carbça
neutral atom minus 4 electrons owned 0.
•
Oxygen has 6 valence elçctrons and I bond. It owns? total valence electrons. The formal
charge =6 valence electrons on the neutral atom minus? electrons owned —1.
‘
•
Oxygen, has the same formal charge as oxygen
=
—1.
Oxygen has 4 valence electrons and 2 bonds. It owns 6 total valence electrons. The formal
2
charge =6 valence electrons on the neutral atom minus 6 electrons owned =0.
The sum of the formal charges, 0 + (1) + (—I) +0= —2, must always equal the charge on the Ion (or
molecule4 if that’s what you are dealing with). Your textbook says that if you can write nonequtvalent Lewis
structures (different numbers of single and double bonds) for a molecule or ion, 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.
I
Example 8.12 A Formal Charges
.
1
Assign formal charges to each atom in the following resonance structures of CO
1.
O=C=O
a
2.
:
b
oc—o:
b
a
Which structure is more iikCly to be correct?
Solution
Lets establish formal charges for each atom.
Structure 1:
Cowris 4 electrons. Formal charge =0
0 owns 6 electrons. Formal charge 0
°b owns 6 electrons. Formal charge =0
Structure 2:
C owns 4 electrons. Formal charge 0
Oa owns 5 electrons. Formal charge +1
—‘1.
°b owns 7 electrons. Formal charge
Structure 1 is more likely because all formal charges are zero.
{ote: Structure 2 can be represented with its formal charges:
oc—o:
+
C)
(b)
L’%)
0
CD
CD
C)
If)
(‘3
CD
CD
(‘3
CD
C)
(I,
I
0
(‘3
CD
CD
CD
7tY
P
0