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Lewis Structures: Notes & Practice
General Process
1. Determine the total number of valence electrons of the atoms in the compound.
2. Arrange the atoms to show bonds between them. The central atom often has the smallest
electronegativity value, and generally appears once in the formula. Remember that hydrogen cannot be
the central atom because it can only form one bond.
3. Distribute the remaining electrons around atoms that do not yet have a full octet. If you have distributed
all of the electrons and all atoms have a full octet, you are done.
4. If you find that there are some left over electrons or one or more atoms do not have an octet then you
will have to create double or triple bonds. Use the electrons from atoms that already have a complete
octet.
For Ions
It is also possible to draw Lewis structures for ions. When drawing for ions, you need to add or subtract the
number of electrons that were gained or lost from ion formation. For example, H3O+ has 9 electrons, but due to
the positive charge, it has lost an electron, and so you only have 8 to configure. OH- has 7 electrons, but due to
the negative charge, it gained an electron, and so you have 8 to configure. Lewis structures for ions are drawn
in brackets with the charge outside on the upper right corner.
Multiple Structures: Calculating Formal Charge
Sometimes there are multiple Lewis structures that satisfy the octet rule. To determine the best Lewis structure
in these situations, we calculate the formal charge of each atom in the molecule. 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.
The rules:
1. The formal charge should be equal to zero for a compound or the charge of the ion for an ion.
2. The best Lewis structure will have atoms with formal charges as close to zero as possible.
3. The best Lewis structure will have any negative charges on the more electronegative atoms.
To calculate formal charge: Total of valence electrons assigned– (0.5(bonding electrons) + lone pair electrons)
1. Draw the possible Lewis structures for a molecule that satisfy the octet rule.
2. Assign electrons to the atom:
a. All unshared or nonbonding electrons are assigned to the atom on which they are found.
b. For any bond—single, double, or triple—half of the bonding electrons are assigned to each atom
in the bond.
3. Subtract the number of electrons assigned to the atom from the number of valence electrons.
Resonance Structures
We sometimes encounter molecules and ions in which the experimentally determined arrangement of atoms is
not adequately described by a single Lewis structure. As a result, we see that the electrons resonate between
two or more arrangements—giving the name resonance structures. Resonance structures show the same
placement of atoms, but different arrangements of electrons. Only Lewis structures with double or triple bonds
can have resonance structures.
Exceptions to the Octet Rule
As you are learning in chemistry, there are often exceptions to the rule. The octet rule is no different. The
exceptions to the octet rule are of three main types:
1. Molecules and polyatomic ions containing an odd number of electrons.
2. Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons.
3. Molecules and polyatomic ions in which an atom has more than octet of valence electrons.
Draw the Lewis structures for the following:
a) Phosphorus tribromide
f) oxygen
b) (Diazine) Dinitrogen dihydride
g) nitrogen
c) Methanol CH3OH
h) formaldehyde CH2O
d) Nitrite ion
i) formate ion (CHO2-1):
e) Ethene C2H4
j) Acetylene C2H2
k) SBF (keep boron central)
p) bromine
l) HBr
q) iodine
m) C2H5OH (ethanol)
r) fluorine
n) dinitrogen tetrafluoride
s) ozone (O3):
o) sulfur hexafluoride
t) carbon dioxide