OUTER
ELECTRONS
VALENCE ELECTRONS
 The
electrons in the outermost energy
levels of the atom are involved in
chemical reactions


They are the ones that bond with other
atoms
These outermost e- are called valence
electrons
 Valence
electrons: an electron in the
outermost energy level of an atom.
LEWIS DOT STRUCTURES
To
show these outer valence
electrons, a system was
created called LEWIS DOT
STRUCTURES
The Lewis dot structures are
based on the s and p orbital of
an atom
HOW TO DRAW LEWIS DOT
STRUCTURES
 To
draw the structure, you need two
things
 The symbol of the element
 The number of valence e Example: Hydrogen
 Hydrogen symbol: H
 The valence e-: hydrogen is 1s1
Thus, hydrogen has 1 valence
electron
HOW TO DRAW LEWIS DOT
STRUCTURES
 Once
you have the symbol and the
valence e- you draw as follows:
 The
H·
symbol is surrounded by the number of
valence e- (represented by dots)
 There can be a maximum of 8 valence
electrons
 Hydrogen is the exception. It only needs 2
LEWIS DOT EXAMPLES
 Berylliuim
 1s22s2
(Be)
 The
outermost valence electrons are
2 (the 2s is further out than the 1s)
 The symbol
.
Be·
LEWIS DOT EXAMPLES (CONT)
Oxygen
(O)
1s22s22p4
The outermost electrons total 6
valence eThe Dot structure:
..
:O·
·
TRY THESE
1.
2.
3.
4.
5.
Boron (B)
Carbon (C)
Nitrogen (N)
Fluorine (F)
Neon (Ne)
SHORT CUT
Examples
on using the
periodic table
HYDROGEN (THE SPECIAL
CASE)
In
order for hydrogen to be stable, it
needs to have the s orbital filled (2 e-)
This can be done by sharing electrons
Example:
H· + H·  H : H
When the electrons are shared, you
symbolize with a (-)
H : H  H - H
HYDROGEN (THE SPECIAL
CASE)
Each
hydrogen atom
shares both electrons to fill
the s orbital
This is very stable
When electrons are shared
it’s called a covalent bond
COVALENT BOND
Covalent
bond: a bond
formed when two or more
valence electrons are
attracted by the positively
charged nuclei of two atoms
and are thus shared between
the two atoms
OCTET RULE
 The
rest of the elements want to have
there s and p orbitals full (very stable
configuration)
 Therefore atoms may donate or
accept e- from other atoms to have a
full outer valence shell
 Since there are 8 e- in the s and p
orbitals, this is called the octet rule
 Octet Rule: tendency of atoms to gain
or lose e- so that their outer s and p
orbitals are filled with 8 electrons
BONDING BETWEEN ELEMENTS
 Consider
the chlorine (Cl) atom
··
:Cl·
··
 When chlorine binds with another chlorine, you get
·· ··
·· ··
:Cl:Cl:  :Cl – Cl:
·· ··
·· ··
 When a pair of electrons are shared they form a
single bond
 The other three pairs around the chlorine atom
are called the unshared pair
ELECTRONS AROUND THE
ATOM
Unshared
pair: a pair of
valence electrons not
involved in bonding with
another atom
Single bond: a covalent bond
in which one pair of electrons
is shared between two atoms
TRY THESE

1.
2.
3.
4.
Draw the Lewis structures
between the following
elements
H+I
I+I
H + Cl
H+F
RULES FOR DRAWING LEWIS
STRUCTURES
 So
far we’ve done some simple Lewis dot
structures, but how do you draw more
complex structures such as CH3I
 Step 1: Find the central atom and
determine the total valence electrons in
the compound.
 Draw each Lewis structure separately
 Add up the total valence e- of each
element
TOTAL NUMBER OF VALENCE
ELECTRONS
 CH3I
 Carbon
(C) has 4 valence e= 4
 3Hydrogen (H) with 1 e= 3
 Iodine (I) has 7 valence e= 7
14electrons
 NOTE:
Carbon has the most number of
open valence e- (the central atom).
RULES FOR DRAWING LEWIS
STRUCTURES
 Step
2: Arrange the atom’s symbols to
show how the atoms bond
 There needs to be 8 electrons around
each element (or 2 in the case of
hydrogen (H))
 Since Carbon (C) has 4 single
electrons, it will be the center atom
 SEE BOARD
RULES FOR DRAWING LEWIS
STRUCTURES
 Step
3: Compare the number of
valence electrons used in the
structure with the number available
in Step 1
 4 single bonds
= 8 electrons
 3 lone pairs
= 6 electrons
14 electrons
RULES FOR DRAWING LEWIS
STRUCTURES
Step
4: Change each pair of
dots that represent a shared
pair of electrons to a single
dash
Step 5: Be sure that all atoms,
with the exception of hydrogen
(H), follow the octet rule.
TRY THESE
1.
2.
3.
4.
CH3I
CH2Cl2
BrF
C2H6
MULTIPLE BONDS

Atoms can share more than 1 pair of
electrons

Consider the following compound:
 Each
O2
O has a total of 6 valence
e Therefore O2 has a total of 12 e How do you surround both O
with an octet of e- . . . Try it!
MULTIPLE BONDS
To
get each atom to have an
octet of electrons, the O must
share 2 pairs of electrons
When compounds share 2 pairs
of electrons, it is called a
double bond
MULTIPLE BONDS

Consider the following compound: N2
 Each
N has a total of 5 valence
e Therefore N2 has a total of 10 e How do you surround both N
with an octet of e- . . . Try it!
MULTIPLE BONDS
 To
get each atom to have an octet
of electrons, the N must share 3
pairs of electrons
 When compounds share 3 pairs of
electrons, it is called a triple bond
 Atoms do not form bigger than a
triple bond
TRY THESE
1.
2.
3.
CO2
C2H4
HCN
LEWIS DOT STRUCTURES - IONS
So
far we have looked at Lewis
structures for normal
elements/molecular
compounds
How do we do Lewis structures
for polyatomic ions?
ANSWER
 Whenever
you have a negative “-” charge, you
add that many electrons
 Whenever you have a positive “+” charge, you
subtract that many electrons
 For example
 SO4-2
 Normally:
 1S
= 6e 4O = 4 x 6e- = 24e Total = 30e-
ADD THE ELECTRONS
 Since


it is SO4-2
You add 2 more electrons to get the -2
charge
This makes your total 32e-
 Now,
when you make the Lewis structure,
you use 32e When you make the structure, you must
put brackets around it
 SEE BOARD FOR EXAMPLE
TRY THESE
2.
ClO4NH4+

ADDITIONAL PRACTICE
1.
RESONANCE STRUCTURES
 Sometimes
the double bond does
not exist between just 2 elements.
 They sometimes form Resonance
structures
 A possible Lewis structure of a
molecule for which more than
one Lewis structure can be written
EXAMPLE
SO2
..
..
..
:S· + :O· + :O ·
·
·
·
S
= 6 valence e2O
= 2 x (6 valence e-)
18 valence eSee board
TRY THESE
1.O3
2.NO3