Honors Chemistry
Unit 6
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
Lewis Dot Structures
VSPER Structures
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We are learning to:
1. Represent compounds with Lewis structures.
2. Apply the VSEPR theory to determine the molecular geometry of a compound.
We are looking for:
1. Draw Lewis structures for compounds and polyatomics based upon octet, formal charge, and
resonance.
2. Determine the molecular geometry of a Lewis structure using VSEPR.
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Name_____________________________________________________
Family Name
Electron configuration
Valance Shell
Electrons
Electron Dot
Notation
Oxidation #
Alkali Metals
Alkaline Earth
Metals
Boron Family
Carbon Family
Nitrogen Family
Oxygen Family
Halogens
Noble Gas
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Introduction to Lewis Structures
1. Draw electron dot for each element
2. Determine if bond will be ionic or covalent
a. Metal + nonmetal  Ionic
b. Nonmetal + nonmetal  Covalent
3. If Ionic: Draw arrows showing the electrons transferring from cation to anion
a. Metals will show a positive charge
b. Nonmetals will show a negative charge
4. If Covalent: Circle the two electrons that make a shared bond. C, N, O, F must obey the octet rule and
achieve 8 valence electrons by covalent bonding (H must always only have 2).
Examples : (also name each compound)
Na2O Ionic or Covalent?
CH4 Ionic or Covalent?
AlCl3
Ionic or Covalent?
NBr3
Ionic or Covalent?
CO
Ionic or Covalent?
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Formal Charges & Lewis Structures
Used when one or more atoms have fewer or more bonds than usual; for example, sometime, C only forms 3 bonds
instead of its usual 4. Also for elements that do NOT have to obey the octet rule (8 valence electrons). C, N, O, F must
obey the octet rule.
On a molecule, formal charges should add up to zero:
:C
O:
On a polyatomic ion, formal charges must add up to the charge on the ion.
[ :Cl O: ] ..
..
..
..
When more than one Lewis structure is possible, use the following rules:
1. A Lewis structure with formal charges of zero is preferable to one with non-zero formal charges. Small formal
charges are preferable to large formal charges.
2. Lewis structures with negative formal charges on the more electronegative atom are more likely than Lewis
structures with negative formal charges on the less electronegative atom.
3. Lewis structures with unlike charges close together are more likely than Lewis structures with opposite charges
widely separated.
4. Lewis structures with like charges on adjacent atoms are very unlikely.
Ex)
A)
:Ö
..
C
O:
B) Ö
..
C Ö
..
..
C) :C.. O O:
The structure with all formal charges equal to zero is the better structure.
Which of the following is the best Lewis structure for formaldehyde, CH2O?
A)
H C Ö
..
H
..
B) C.. O H
..
C) H C Ö H
H
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Resonance Structures:
Even after using the rules for formal charges, sometimes more than one Lewis structure can be written. This often
occurs when using unshared electron pairs to form multiple bonds. In this case both (or more) must be written. These
are called resonance structures. In reality, the resulting structure is a hybrid of all the resonance structures. Use a
double headed arrow
between them.
*Very important*: when writing resonance structures, all the atoms must be shown in the
same positions; only the positions of the electrons are different.
We can know the position of the atoms by experiment, but the position of the electrons cannot be determined; only the
probability of finding an electron in a certain region can be known.
Ex) thiocyanate ion, SCN..
..
..
S.. C N..
But not
:S.. C N:
..
:S
.. N C:
Why?
Nitrate, NO3..
..
..
..
..
..
..
..
..
..
..
..
:O N O:
O N O:
:O N O
:O:
:O:
:O:
..
..
Acetate, C2H3O2-
H
..
H C C=O
..
H :O:
..
..
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Lewis Structures
If single bonds are used and not all atoms have an octet (Except IA, IIA, and IIIA) – Try unshared pairs
then double and triple bonds:
CO2
O
C
O
CCl2O
Cl
C
O
Cl
C2Cl2
Cl
C
C
Cl
1. O2
4. HClO2
2. NCCN
5. N2H4
3. O3
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Summary of Writing Covalent Lewis Structures
1. Using whatever information available, write the symbols for the elements in the correct
arrangement.
2. Calculate the total number of valence electrons. Remember to adjust amount for
polyatomic ions!
3. Place one pair of electrons (single line) between each pair of bonded atoms.
4. Beginning at the outside of the formula, place the remaining electrons in pairs until
there are eight electrons around each atom (two for hydrogen) or all electrons have
been used.
a. If there are extra electrons, place them on the central atom.
b. Atoms in the 3rd period and beyond can have more than eight electrons around
them and can form more than 4 bonds, but elements in the 2nd period cannot.
5. If not enough electrons are available to give all atoms (except H) an octet, move
unshared pairs to form double or triple bonds.
a. However, Be, B, and other Group IIA elements may have fewer than eight
electrons.
6. Examine your Lewis structure to see if resonance structures are needed. (Is more than
one position possible for multiple bonds)
7. Check the formal charges of the atoms. It should be at the lowest possible, even if the
octet rule is not fulfilled (exception: C, N, O, F must have an octet and H two electrons).
Want to try to have a formal charge of zero for the central atom.
8. Check your answer. Does the Lewis structure show:
a. the correct number of atoms?
b. the correct number of electrons?
c. the right number of electrons around each atom?
d. the minimum number of formal charges?
e. The brackets around and the charge on it if it is a polyatomic ion?
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Lewis Structures
Formula
(Give the name of the
compound under the
formula)
Ionic
Or
Covalent?
Number
Of
Valence
Electrons
Lewis
Structure
1. Rb3N
2. NF3
3. SCl2
4. InBr3
5. PCl5
6. SF6
7. SeCl4
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Formula
**(You do NOT have to
name these)**
Ionic
Or
Covalent?
Number
Of
Valence
Electrons
Lewis
Structure
8. C2H6O
(CarbonCarbon-Oxygen
bonding
pattern)
9. CCl2F2
(Carbon in
middle)
10. NH2Cl
11. SOCl2
(Sulfur in the
middle)
12. CH3COCH3
C-C-C
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Lewis Structures for
Polyatomic Species
**Also, Name Each Ion!**
Examples: NH4+
and
PO43-
When the charge is +
Subtract the charge value from the total number of valence electrons
When losing electrons – they usually come off the central atom!!
When the charge is –
Add the charge value to the total number of valence electrons
When gaining electrons - usually don’t add to the central atom!!
Example: NH4+ _____________________________
Atom
How many
x
# of e-s Total for each element
N
1
5
=
5
H
4
1
=
4
1+
losing 1
=
-1
Total valence electrons = 8
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Example: PO43- ____________________________________
Atom
How many
x
# of e-s Total # of electrons
P
1
5
=
5
O
4
6
=
24
gaining 3
=
3
3-
Total valence electrons = 32
Example: BrO3
Atom
-
______________________________________________________________
How many
x
# of e-s total # of electrons
Br
1
7
=
7
O
3
6
=
18
1-
gaining 1
=
1
Total valence electrons = 26
CO32- _________________________________________
Atom
How many
x
# of e-s total # of electrons
C
1
4
=
4
O
3
6
=
18
2-
gaining 2
=
2
Total valence electrons = 24
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Show math calculations and then draw the Lewis structure
1.
H3O+ ____________________________________
2.
NO3- _________________________________________
3.
CN-
4.
SO32- ________________________________________
5.
OH- __________________________________________
_cyanide____
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VSEPR Theory
(Valence Shell Electron Pair Repulsion)
 Predicts the molecular shape of the resulting molecule
 Electrons on central atom arrange themselves as far apart as possible
 Unshared pairs on the central atom repel the most
 Shared pairs on the central atom repel the least
 To get the shape ONLY LOOK AT WHAT IS CONNECTED TO THE
CENTRAL ATOM!!!!!
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VSEPR Theory
Type
Overall Geometry
AX2
AX3
Bond Angle (o)
Examples
Linear
180
SiS2, CO2
Trigonal Planar
120
BCl3, H2CO
Bent
<120
SO2
Tetrahedral
109.5
CH4, (SO4)-2
Trigonal Pyramidal
<109.5
NH3
AX2E
AX4
Tetrahedral
AX3E
Molecular Geometry
(107)
AX2E2
Bent
<<109.5
H2O
(104.5)
AX5
AX4E
Trigonal
Bipyramidal
Trigonal
Bipyramidal
See-Saw
180
PCl5
120
<180
<120
(177)
SF4
(104)
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AX3E2
T-Shape
<90
(87.5)
ClF3
AX2E3
Linear
180
XeF2, I3-1
Octahedral
90
SF6
AX5E
Square Pyramidal
<90
IF5
AX4E2
Square Planar
90
(BrF4)-1, XeF4
AX6
Octahedral
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Here are some rules about using VSEPR to describe molecular geometries:
1. Electron pair repulsions decrease in the following order:
a. (nonbonding-nonbonding)> (nonbonding-bonding)> (bonding-bonding)
b. triple bond > double bond > single bond
2. Nonbonding pairs occupy equatorial positions in the trigonal bipyramid.
3. Two nonbonding pairs in an overall octahedral structure occupy positions across from one another.
4. Bond angles decrease with increasing electronegativity of the non-central atoms, provided that there is at least one
nonbonding pair in the species. For example, the F-N-F bond angle in NF3 is smaller than the H-N-H bond angle in
ammonia. Picture the electrons being closer to the more electronegative element fluorine than to the nitrogen. This
allows the nonbonding pair to squeeze the bonding pairs closer together, resulting in a smaller bond angle. In ammonia,
the bonding electrons are closer to the more electronegative element N so the electrons are already close to each other.
The bond angle doesn't get any smaller.
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VSEPR Practice (molecular geometry)
1) Linear
a. linear= 0 lone pair
Draw:
SiS2
CO2
2) Trigonal Planar
a. trigonal planar= O lone pair
Draw:
BCl3
H2CO
b. bent= 1 lone pair
Draw:
SO2
O3
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3) Tetrahedral
a. Tetrahedral= 0 lone pair
Draw:
CH4
(SO4)2-
b. trigonal pyramidal= 1 lone pair
Draw:
NH3
c. bent= 2 lone pairs
Draw:
H2O
<<109
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4) Trigonal bipyramidal
a. Trigonal bipyramidal= 0 lone pairs
Draw:
PCl5
b. See-saw= 1 lone pair
Draw:
SF4
c. T-shaped= 2 lone pairs
Draw:
ClF3
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d. linear= 3 lone pairs
Draw:
XeF2
I3 -
5) Octahedral
a. octahedral= 0 lone pair
Draw:
SF6
b. square pyramidal= 1 lone pair
Draw:
IF5
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c. square planar= 2 lone pairs
Draw:
(BrF4)-
XeF4
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Lewis Structure/VSEPR Problems
Formula
(write the name
under the
formula)
Lewis
Structure
Number of
Bonded Atoms
Number of
Lone Electron
Pairs
Molecular
Geometry
(name &
sketch)
1) SiF4
2) BBr3
3) NF3
4) H2O
5) AsCl5
6) SF6
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Lewis Structures and VSEPR Practice
**Only do the VSEPR Shape on Covalent Structures**
Formula
(write the name
under the formula)
Ionic or
Covalent
Lewis Structure
VSEPR Shape
Drawing
Shape Name
1) KCl
2) BH3
3) CCl4
4) CO32-
5) SF2
6) Na2O
7) AsBr3
8) H3O+
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Name:_________________________________
Formula
Compound Name
Lewis Structure/VSEPR Practice WS
Lewis
Structure
Resonance?
Yes or No
VSEPR Structure (include at
least 1 angle)
VSEPR
Name
1) SiI4
2) Al N
3) NO3-
4) BeCl2
5) PF3
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Formula
Compound Name
Lewis
Structure
Resonance?
Yes or No
VSEPR Structure (include at
least 1 angle)
VSEPR
Name
6) TeH6
7) Rb2O
8) PI5
9) PO4 3-
10) XeF2
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