This document provides instruction on drawing Lewis structures to represent covalent bonding. It introduces Lewis structures and notes that they show valence electrons as either shared pairs forming bonds or lone pairs not involved in bonding. It describes the octet rule where atoms gain, lose or share electrons to achieve eight valence electrons, and exceptions for hydrogen and boron. Rules are given for drawing Lewis structures, including determining atoms and electrons, arranging atoms, adding lone pairs to achieve octets, and counting electrons. Examples are provided for practice. Multiple bonds between carbon, nitrogen and oxygen are also discussed.

1. LESSON 3: DRAWING LEWIS
STRUCTURES
Covalent Bonding Unit

2. TARGETS
 I can draw Lewis Structures for covalent
compounds.

3. DEFINITION
 Lewis Structures – visual representation of
covalent bonding that indicates where the
valence shell electrons are in the molecule.
 Shared electron pairs are shown as lines and
lone pairs (pairs of electrons not involved in
bonding) are shown as dots

4. OCTET RULE
 Atoms tend to gain, lose, or share electrons
until they are surrounded by eight valence
electrons (octet). An octet consists of full s and p
orbitals.

5. EXCEPTIONS TO THE OCTET RULE
 Duet Rule - hydrogen only has 2 valence
electrons after bonding
 Boron only has 6 valence electrons after bonding

6. LEWIS STRUCTURE RULES
1. Determine the type and number of atoms in the
molecule.
2. Determine the total number of valence electrons
available in the atoms to be combined.
3. Arrange the atoms to form a skeleton structure for the
molecule. If carbon is present, it is the central atom.
Otherwise, the least electronegative atoms is central
(except for hydrogen which is never central). Then
connect the atoms by electron-pair bonds.
4. Add unshared pairs of electrons to each nonmetal atom
(except hydrogen) such that each is surrounded by eight
electrons.
5. Count the electrons in the structure to be sure that the
number of valence electrons used equals the number
available. Be sure the central atom and other atoms
besides hydrogen have an octet.

7. PRACTICE
 Draw Lewis Structures for the following:
a. CH3I
b. NH3
c. H2S
d. PF3
e. IBr
f. F2O

9. MULTIPLE COVALENT BONDS
 In writing Lewis structures for molecules that
contain carbon, nitrogen, or oxygen, one must
remember that multiple bonds between pairs of
these atoms are possible.
 A hydrogen atom has only one electron and
therefore always forms a single covalent bond.
 The need for a multiple bond becomes obvious if
there are not enough valence electrons to
complete octets by adding unshared pairs.

10. LEWIS STRUCTURE RULES
1. Determine the type and number of atoms in the molecule.
2. Determine the total number of valence electrons available in the
atoms to be combined.
3. Arrange the atoms to form a skeleton structure for the molecule.
If carbon is present, it is the central atom. Otherwise, the least
electronegative atoms is central (except for hydrogen which is
never central). Then connect the atoms by electron-pair bonds.
4. Add unshared pairs of electrons to each nonmetal atom (except
hydrogen) such that each is surrounded by eight electrons.
5. Count the electrons in the structure to be sure that the number of
valence electrons used equals the number available.
6. If too many electrons have been used, subtract one or more
lone pairs until the total number of valence electrons is
correct. Then move one or more lone electron pairs to
existing bonds between non-hydrogen atoms until the
outer shells of all atoms are completely filled.

11. PRACTICE
 Draw Lewis structures for
a. CH2O
b. CO2
c. HCN
d. C2H2