Ch. 6.2 Notes
DRAWING AND NAMING MOLECULES
Lewis Dot Structures
 Represent elements and their valence e Valence e- are in outermost s & p
 e- are represented by dots placed on the 4 sides of
the element’s symbol
 Put 1 e- on each side, then begin pairing
 Max. of 8 e-/dots  follows octet rule
H•
Li•
•
Be•
•
•B•
•
•C •
•
•• • •
N
•
•• • •
O
••
••
••F •
••
• •• •
• Ne •
••
Representing Bonds
 H2 
 Cl2 
H•
• •••
• Cl
••
•H
•• •
• Cl•
••


H—H
••
••
•• Cl—Cl••
••
••
 e- that are not involved in the bond are unshared
pairs  called lone pairs
 When one pair of e- is shared, represented by a
single line (—)

Called a single bond
 HCl 
H•
••
• Cl••
••

•• •
•
H—Cl
••
Drawing Lewis Dot Structures - Steps
 1. Draw Lewis structure for each atom in compound.
Place one e- on each side before pairing. Count total
# of valence e 2. Arrange the Lewis structure to show how atoms
bond in the molecule



Halogens and hydrogen are usually at the end
Carbon is usually in the middle
…Or atom with lowest electronegativity is placed in the middle
Drawing Lewis Dot Structures - Steps
 3. Distribute the dots (e-) so that each atom is getting
its octet
 4. Draw the bonds that form from shared e 5. Check your work by making sure your final Lewis
dot structure represents the same amount of e- you
started with
Example
 Draw Lewis structure for CH3I
•
• C•
•
H•
H•
 Arrange the atoms
H•
••
• I ••
••
H
H C I
H
 Distribute the dots (e-)
H
•• ••
H• • C• • •I ••
••
•
H
Example
 Distribute the dots (e-)
H•
•
••
• • • • ••
H C• •I•
•
H
 Draw the bonds
••
H
•
•
H—C—I
••
H
 Verify the structure
Practice
 H2S
 CH2Cl2
 NH3
 C2H6
 CH3OH
 Challenge!
Lewis Structures for Polyatomic Ions
 1. Draw Lewis structure for each atom in compound.
Place one e- on each side before pairing. Count your
total valence e-. You will then add or subtract ebased on your charge.

Ex. NH4+ — 9 valence e-, has +1 charge — 9-1 = 8 valence e-
 2. Arrange the atoms, distribute your dots
(remember to use your adjusted total)
 3. Draw the bonds, verify you have the correct # e 4. Put brackets around the ion, write the charge as a
superscript outside the bracket
Example
 Draw Lewis structure for SO42•
•• S •
••
•
•• O •
••
•
•• O •
••
•
•• O •
••
•
•• O •
••
# e- = 30 + 2 = 32
 Arrange atoms/distribute dots
••
•• O ••
•• ••
••
•• O • • S • • O ••
•• ••
••
•• O •
•• •
 Draw bonds/verify structure
•• •• ••
2O
••
•••
•O—S—O••
••
••
• •
•O
•• •
Practice
 ClO3 H3O+
Classwork/Homework
 Draw the Lewis structures for the following
compounds:







H2O
CH4
PBr3
N2H4
CCl4
NI3
ClF4+
Multiple Bonds
 Atoms can share more than 1 pair of e- in a covalent
bond

Sometimes just sharing 1 pair of e- wouldn’t fill the octet
 Example: O2



••
•• O •
•
••
•• O•
•
Each O only has 6 valence eSharing 1 pair would still leave an unpaired e- on each
Sharing 2 pairs would give each atom its octet
••
•O
•
••
= O ••
Practice!
 C2H4 has a double bond. Draw its Lewis structure.
Triple Bond!
 N2 has a triple bond (3 pairs of e- shared). Try
drawing its Lewis structure.
More Practice with Multiple Bonds
 CH2O
 CO2
 CO
 C2H2
 HCN
Resonance Structures
 When a molecule has 2 or more possible Lewis
structures
 Example: O3
• • •• • •
•• O—O=O ••
••
••
•• ••
••O=O—O
••
••
Relative Lengths of Bonds
 Bond length also determined by type of bond
 Single > Double > Triple
Naming Binary Covalent Compounds
 Name first element first
 Then name second element—change ending to –ide
 Prefixes used to indicate how many of each element
is present

Only exception is when you have 1 of your first element, no
prefix
Naming Binary Covalent Compounds










MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
1
2
3
4
5
6
7
8
9
10
 If adding prefix to name that starts with a vowel, don’t include the
“a” in the name

5 oxygens would be pentoxide, not pentaoxide
Examples
 CO
 Carbon monoxide
 S2F10
 Disulfur decafluoride
Your turn
 SbCl5
 N2O
 SCl4
 P4O6
 SO3
 SiO2