Lesson 38 Lewis Dot
Diagrams and Structures
Objectives:
The student will draw Lewis Dot
Diagrams for various elements.
The student will construct Lewis
Structures for simple compounds.
The student will use Lewis Structures to
explain double and triple bonds, and
resonance structures.
PA Science and Technology Standards: 3.1.10.B
I.
Lewis Structures
a.
Electronegativity differences don’t tell us
everything we need to know.
b.
We can use dots to represent valence
electrons involved in bonding.
c. Valence electrons – electrons present in
the outermost energy level of an atom.
d.
Simplest example – hydrogen
i.
One hydrogen – H.
ii.
Two together – H:H
Atoms are represented with the number of dots
around them that is the same as their group number,
if they are in an A group.
f. Examples of single elements:
g.
If they are in a B group (transition metals), they will have a
number of electrons determined by the charge they carry.
h.
Follow the following rules to determine how to make a
Lewis structure:
i.
Determine the total number of valence electrons in
the compound
ii.
Arrange the atoms’ symbols to show how they are
bonded and show valence electrons as dots
iii.
Compare the number of valence electrons used in
the structure to the number available from step one
iv.
Change to a single dash each pair of dots that
represents two shared electrons.
v. Be sure that all atoms, with the exception of
hydrogen, follow the octet rule.
vi. General rules to follow:
1.
Hydrogen and halogen atoms usually
bond to only one other atom in a
molecule, and are usually on the outside
or end of a molecule.
2.
The atom with the smallest
electronegativity is often the central atom.
3.
When a molecule contains more atoms of
one element than the others, these atoms
often surround the central atom.
vii.
Two atoms can share more than one electron
pair – try doubling up pairs of electrons if you
need more electrons than the original
calculations allow to satisfy the octet rule.
viii. Examples:
i.
A Single bond is the sharing of one pair of
electrons between two atoms.
j.
A Lewis structure is a diagram showing the
arrangement of valence electrons among the
atoms in a molecule.
II.
Two atoms can share more than one electron pair
a.
Carbon is a good example of an atom that
can share more than one pair of electrons -it can share two
or three with another carbon atom
b.
Several other elements, such as nitrogen,
oxygen, and sulfur, can also share more
than one pair.
i. A double bond is a covalent bond
formed by the sharing of two pairs of
electrons between two atoms
ii.
A triple bond is a covalent bond
formed by the sharing of three pairs of
electrons between two atoms.
c.
If no arrangement of single bonds
provides an appropriate Lewis structure,
consider looking into forming multiple bonds.
d.
a. Examples:
Examples:
Lewis Structures Virtual Lab 2
Video
III.
Sometimes no single Lewis structure is correct
a.
For some molecules, it is not possible to draw one
correct Lewis structure.
b.
NO2 is a good example of this. Example:
c.
Things to think about
i.
Atoms with unpaired electrons are somewhat
unstable compared to atoms with all electrons
paired, but they are possible.
ii.
Situations where an octet does not form are
normally unstable, but possible in some cases
also.
iii.
Before such a rule breaking structure can be
considered accurate, either an octet must be
impossible, or a non-octet structure must be found
to be experimentally more stable.
iv.
Example
d.
Often several equivalent structures can be drawn – These
are called resonance hybrids.
i.
Originally, it was thought that the structures
resonated between the possible structures
ii.
It was later determined that the actual structures
are an average, or combination, of these
structures.
iii.
Examples:
IV.
Lewis structures can be drawn for polyatomic ions.
a.
Polyatomic ions are held together by covalent bonds
– we can draw Lewis structures for them just as
easily.
b.
Ammonium – NH4
i. The ion is prepared from ammonia, NH3, and a
proton (the nucleus of a hydrogen atom)
ii. The hydrogen shares a pair of electrons donated
by the nitrogen
iii. This produces a coordinate covalent bond – a
bond where two atoms share a pair of electrons
donated by only one of the atoms.
iv.
Once it is formed, this coordinate covalent bond
is indistinguishable from a normal covalent bond.
v.
Brackets are placed around the structure, with
the charge on the ion shown outside the
brackets.
When you count up electrons that should be in
the structure, you need to account for the charge
on the ion.
vi.
1.
If it is a positive charge, subtract that
number of electrons from the total.
2.
If it is a negative charge, add that number
of electrons to the total.
vii.
Examples:
Questions:
1.
Create Lewis dot diagrams for each of the following
elements:
a.
Na
He
S
P
Mg
I
O
Al
Ne
b.
c.
d.
e.
f.
g.
h.
i.
1.
Create Lewis dot diagrams for each of the following
compounds:
a.
HI
O2
CCl4
SO2
ICl
Cl2
H 2O
CH3COOH
NO3SO4-2
PO4-3
N2
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
 CO32 NBr3
 SiO2
 BI3
 ArO2
 AlBr3
HCP
O2
P2
HCBrF2
CS2
GeI4