Ch. 6
Lewis Structures and
Molecular Geometry
How to draw Lewis Structures, or dot
diagrams of e- bonding
 Only the Valence e- are involved
 Most atoms strive to get 8 e- in their outer
principal electron layer. This is the “Octet
Rule.”

Here’s how many valence e- each
group has. Write them in on your chart!
In a Lewis Structure
Each dot represents
1 electron.
 Each line represents
2 shared electrons in
a covalent bond.
 Sometimes the lines
are shown as dots;
see this diagram:

How to draw Lewis Structure for
the OCl- (bleach) ion.
1. Count the valence e- available.
 Add electrons for negative charges.
 Subtract electrons for positive charges.

Step 2:

Draw a “skeleton” structure using single
bonds (lines) for pairs of shared electrons.
Step 3

Distribute the rest of the e- around the
atoms so each atom has 8 electrons.
There are 3 things that can occur with
your electrons on the central atom:



1. There are just enough e- to go around. Every
atom has 8 and H atoms have 2.
2. If there aren’t enough e- to go around, then
move a pair of e- next to an existing single bond
to make a double or triple bond.
3. If there are too many e-, then the atom may
have an “Expanded Octet.” Put the extra epairs on the central atom.
An example of an “expanded
octet”

Here’s the Lewis Structure for XeF4
Important Tips:
Carbon is usually a central atom, while H,
O, and the Halogens are usually terminal
atoms with ONLY one single bond.
 The FIRST atom listed in the chemical
formula is usually the central atom.
 Carbon always has 4 lines (4 covalent
bonds) coming off of it.

Draw the Lewis Structure for
CH3OH, methanol.
Draw the Lewis Structure for
the sulfite ion, SO32-
Draw the Lewis Structure for
the Nitrate ion, NO3 .

Stuck? Look at the “Important Tips” we
just covered!
Exceptions to the Octet Rule:
Some light elements don’t have 8 e- on
the central atom.
 Examples: H, Li, Be, and B
 Lewis Structures for BeF2 and BF3

Part 2: Predicting Molecular
Geometry (Shapes)
Molecular shapes can be predicted using
the Valence Shell Electron Pair Repulsion,
VSEPR principle.
 All e- pairs try to get as far away from
each other as possible.
 You must draw the Lewis structure before
you can predict the geometry.

A good
reference
chart:
How to predict the “Repulsion
Angles” in molecules:
Add up the number of “forces” pushing
away from the central atom.
 Each lone e- pair and atom(s) bonded to
the central atom, count as a “force.”
 Double or Triple bonds behave like just
one “force.”

The 3 basic “Repulsion Angles”
2 Forces: make a Linear
molecule with 180 ° angles
off the central atom.
 3 Forces: repel in a
Trigonal Planar pattern with
120 ° angles.
 4 Forces: repel in a
Tetrahedral pattern with
109.5 ° angles.

The names of the possible
shapes.
I presented only 3 possible repulsion
angles, but 5 possible molecular shapes
exist given these 3 angles.
 The molecule’s shape is named after the
3D arrangement of the atoms.
 Unshared e- pairs (dots) on the central
atom influence the shape but are not seen.

How do you tell if a molecule is
“polar?”

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A molecule is “polar” when it has more electrons
on one side than on the other. This means it
has a negative side with more electrons and a
positive side with less electrons.
Generally, if a molecule is perfectly symmetrical
with no unshared e- pairs, then it is non-polar.
Unshared pairs usually make it polar.
If the terminal atoms are different, then it's polar
(different charge on each end)
What is the molecular
geometry of BF3 ?

120 ° angles, Trigonal Planar, non polar
What is the geometry of
a water molecule?

105 ° angles, Bent, very polar.
Polar or Nonpolar?