Lewis Dot Diagrams
• A study of chemical bonding would not be complete –
even in these brief notes – without teaching about
Lewis Dot Diagrams (Standard 2e). Professor Gilbert
N. Lewis spent much of his career at UC Berkeley,
teaching chemistry. He was an accomplished
scientist and mentored at least one Nobel Prize
laureate. Lewis himself was nominated for the prize
many times but appears to have been passed over for
political or non-science reasons. In 1916 he developed
the notation called “dot structures” or “dot diagrams”
as a teaching aid to help students understand atomic
structure and bonding. He discovered the covalent
bond.
 Dot structures allow students and chemists alike
to understand the way molecules are assembled
from atoms. It is based on the valence bond
theory and is a logical tool. By use of dot
structures, we are able to understand that there
can be more than eight electrons in a valence
orbital, just not in periods 2 and 3.
 First I’ll explain the method, and then we will
look at some examples. Next you will have the
chance to practice drawing structures yourself.
 In a two-atom system, the dot diagram is no
challenge. It is drawn with available electrons
around and shared by the atoms.
 With 3 or more atoms, the distribution becomes a
little trickier. Specific rules are used.
1. The least electronegative atom goes in the
center of the diagram.
2. Hydrogen atoms are usually at the outside.
3. The total number of valence electrons are
counted. If other electrons have been gained or
lost, they are added or subtracted, respectively.
4. Bonds are drawn from the center atom to the
others, usually, not always based on the octet
rule.
5. Other electrons are “dotted” in to the diagram
starting from the outside atoms.
6. To verify that the structure is the best that it
can be, use the Formal Charge equation:
F.C. = V.E. – L.P. – B.P./2
This means the Formal Charge on the atom is
equal to the total number of its valence electrons
(V.E.) minus the lone pair electrons (L.P.) minus
the number of bound pair electrons (B.P.) divided
by two. Strive for a value of zero or as close to
zero as possible.
• Here are some examples:
 Show the Lewis structure for SCl2.
S = sulfur = 6 valence electrons
Cl = chlorine = 7 valence electrons. There are
two chlorine atoms, so the total electrons available
for bonding is 6 + 7 + 7 = 20. Sulfur is less
electronegative than chlorine. Write the structure
like this:
Cl – S – Cl
The lines (--) indicate a single bond, using two
electrons each, leaving 16 to place:
.. . . ..
:Cl – S – Cl:
..
..
..
Analyze: Formal charge for chlorine is
7 VE – 6 LP – 2(BP)/2 = 0
This accounts for all electrons. The molecule’s
shape looks like a deformed pyramid because of
the unbonded electrons on the sulfur (in the
center of the pyramid):
 Show the Lewis dot structure for phosphorus
pentachloride, PCl5.
P = phosphorus = 5 valence electrons
Cl = chlorine = 7 valence electrons
There are 5 chlorine so there are 35 + 5 = 40
electrons.
Phosphorus is less
electronegative than
chlorine . Write the
structure like this:
Cl
|
Cl – P - Cl
/ \
Cl Cl
The lines indicate single bonds and there are five
of them so this uses 10 electrons leaving 30 more
to place. Start at the outside (chlorine) and place
them:
..
:Cl:
| ..
:Cl – P – Cl:
..
..
..
.. / \ ..
:Cl: :Cl:
..
..
Notice that phosphorus has 5 bonds, and it can
because it is in the third row of the periodic table,
where the octet rule does not need to apply. The
shape of the molecule is a double pyramid. You
can use the formal charge equation to check your
work .
 Show the Lewis structure for CH3CO2H
(acetate ion. Carbon contributes 4
|
••
electrons, and there are 2 carbon atoms.
H–C–C=O
Hydrogen has one electron each,
••
1| |
and there are 3, and there are two
H :O:
oxygen atoms, with 6 electrons. Also,
••
notice there is one free electron for a
•
total of 24
electrons. Carbon is least
electronegative. The carbon atoms go
in the
middle:
Note that the brackets and the (1-) indicates the extra
electron. Why is the single-bonded oxygen allowed?
Because it has the extra electron and is an anion, waiting
for its cation partner.
• Now it is time for you to practice: Here are
sample problems. Use the periodic table to find the
number of valence electrons. If you get into trouble,
there are a couple of very good tutorials online1.
1
http://www.chem.ucla.edu/harding/lewisdots.html
http://www.kentchemistry.com/links/bonding/lewisdotstruct.htm
1. CH3CH2OH
2. H2SO4
3. H3PO4
4. NH2OH
5.
6.
7.
8.
NO2
OHCO
BeH2
9. BH3
10. CH4
11. NH3
12. BH4-
13. NH4+
14. H3O+
15. BCl3
16. SiF62-