

Recall: atoms want noble gas configurations
Octet Rule: atoms will gain, lose, or share
electrons to achieve a noble gas configuration
 Typically 8 valence electrons

Atoms will bond with each other to achieve a
full octet
F

F
A pair of shared electrons can be represented
by either with 2 dots or with a dash
F

F
Unshared electrons are called lone pairs

Types of bonds:
 Single bonds: bond containing only 2 electrons
 Multiple bonds: bond containing more than 2
electrons
▪ Double bond: 4 electrons (or 2 pairs of electrons)
▪ Triple bond: 6 electrons (or 3 pairs of electrons)

In a particular pair of elements
 Triple bonds are the shortest
 Double bonds are in the middle
 Single bonds are the longest
Bond energy is the
energy required to
BREAK bonds
between atoms

Electronegativity is a periodic trend
 Ability of an atom to attract electrons to itself
when bonded to another atom
 Quantified by the Pauling Scale

Let’s consider three molecules:
 H2, HF, NaF

Take the difference in electronegativities of
two atoms bonded together
 If the difference is 0.5 or lower, the bond is
nonpolar covalent
 If the difference is between 0.5 and 2.0, the bond
is polar covalent
 If the difference is greater than 2.0, the bond is
ionic
The bond in ClF
(chlorine and fluorine)
 The bond in CsBr
 The carbon-carbon
double bond in C2H4


In which of the
following molecules
are the bonds most
polar: H2O, BCl3, PCl5

1) Determine the central atom and place
terminal (“outside”) atoms around central atom
 Central atom typically is the least electronegative
element in compound, the element with only 1
atom, and/or the element written first in compound

2) Count total # of v.e.

3) Bond all terminal atoms to central using
single bond
 Each bond is 2 electrons; subtract from total # of v.e.

4) Complete the octets of terminal atoms w/
remaining v.e.

5) If any electrons left over, put on central
atom

6) Use multiple bonds to complete octet of
any elements where necessary





CH4
H2O
O2
CO2
CN-

Draw the Lewis Dot Structures for the
following:

CS2
NF3
ClO3-



Another way of keeping track of electrons in
a molecule

Formal Charge = (# of v.e.) – (# of associated
electrons)

Ex: Ozone (O3)
Now you try: NO3-



Formal Charge can help us determine the
best Lewis Structure when there are options
Consider the following two skeletal structures
for CH2O. Which one is preferred?

Lewis structures where all formal charges are
zero is preferred

Small formal charges (0 and +/-1) are
preferred to big formal charges (+/-2, +/-3,
etc.)

The best arrangements are where the more
electronegative atoms have the more
negative formal charge

Draw the Lewis Structures for the following
compounds and determine the formal charge
on EACH atom

SO32-

CO32-

Consider the molecule NO3- and its Lewis
Structure

Central atom has fewer than 8 v.e. due to
electron shortage
 Ex: Boron (happy w/ 6); Beryllium (happy w/ 2)

Central atom has fewer than 8 v.e. due to odd
# of electrons (known as radicals)
 Ex: Nitrogen (NO2)

Central atom has more than 8 v.e.
 Ex: Sulfur (SF6) and Xenon (XeF4)
 See pp. 201—204 for examples/explanations

Consider the polyatomic ion: SO42-. What
would be the BEST Lewis dot structure?

What about PO43-?

Draw the Lewis Structure for antimony
pentafluoride (SbF5)

Draw the Lewis Structure for Borane (BH3)

Draw the Lewis Structure for Nitrogen
Disulfide (NS2)