Honors Chemistry
Name: _______________________________________ Date: ___________________ Mods: __________
Ch. 8 – Chemical Bonding, Electronegativity
Differences, and Lewis Structures
A) Summary of Chemical Bonds and Their Properties

Ionic Bond: a non-metallic _____________ TAKES electrons from a metallic _____________ due to a
__________________ difference in electronegativity.
Example:

Polar Covalent Bond: an __________________ difference in electronegativity leads to an UNEQUAL
SHARING of electrons between two ___________________.
Example:

Nonpolar Covalent Bond: two ____________________ with ________________ electronegativity or
only a ______________ difference in electronegativity EQUALLY SHARE electrons.
Example:

Metallic Bond: results from the ATTRACTION between ___________ atoms and the surrounding sea of
__________________.
Example:
1
H
2.1
3
Li
1.0
11
Na
1.0
19
K
0.9
37
Rb
0.9
55
Cs
0.8
4
Be
1.5
12
Mg
1.2
20
Ca
1.0
38
Sr
1.0
56
Ba
1.0
21
Sc
1.3
39
Y
1.2
57
La
1.1
22
Ti
1.4
40
Zr
1.3
72
Hf
1.3
23
V
1.5
41
Nb
1.5
73
Ta
1.4
24
Cr
1.6
42
Mo
1.6
74
W
1.5
25
Mn
1.6
43
Tc
1.7
75
Re
1.7
26
Fe
1.7
44
Ru
1.8
76
Os
1.9
27
Co
1.7
45
Rh
1.8
77
Ir
1.9
28
Ni
1.8
46
Pd
1.8
78
Pt
1.8
29
Cu
1.8
47
Ag
1.6
79
Au
1.9
30
Zn
1.6
48
Cd
1.6
80
Hg
1.7
5
B
2.0
13
Al
1.5
31
Ga
1.7
49
In
1.6
81
Tl
1.6
6
C
2.5
14
Si
1.8
32
Ge
1.9
50
Sn
1.8
82
Pb
1.7
7
N
3.0
15
P
2.1
33
As
2.1
51
Sb
1.9
83
Bi
1.8
8
O
3.5
16
S
2.5
34
Se
2.4
52
Te
2.1
84
Po
1.9
9
F
4.0
17
Cl
3.0
35
Br
2.8
53
I
2.5
85
At
2.1
87
Fr
0.8
88
Ra
1.0
89
Ac
1.1
B) Using Electronegativity to Predict the Polarity of a Chemical Bond
1
H
2.1
3
Li
1.0
11
Na
1.0
19
K
0.9
37
Rb
0.9
55
Cs
0.8
87
Fr
0.8
1.
4
Be
1.5
12
Mg
1.2
20
Ca
1.0
38
Sr
1.0
56
Ba
1.0
88
Ra
1.0
21
Sc
1.3
39
Y
1.2
57
La
1.1
89
Ac
1.1
22
Ti
1.4
40
Zr
1.3
72
Hf
1.3
23
V
1.5
41
Nb
1.5
73
Ta
1.4
24
Cr
1.6
42
Mo
1.6
74
W
1.5
25
Mn
1.6
43
Tc
1.7
75
Re
1.7
26
Fe
1.7
44
Ru
1.8
76
Os
1.9
27
Co
1.7
45
Rh
1.8
77
Ir
1.9
28
Ni
1.8
46
Pd
1.8
78
Pt
1.8
29
Cu
1.8
47
Ag
1.6
79
Au
1.9
30
Zn
1.6
48
Cd
1.6
80
Hg
1.7
5
B
2.0
13
Al
1.5
31
Ga
1.7
49
In
1.6
81
Tl
1.6
6
C
2.5
14
Si
1.8
32
Ge
1.9
50
Sn
1.8
82
Pb
1.7
7
N
3.0
15
P
2.1
33
As
2.1
51
Sb
1.9
83
Bi
1.8
8
O
3.5
16
S
2.5
34
Se
2.4
52
Te
2.1
84
Po
1.9
9
F
4.0
17
Cl
3.0
35
Br
2.8
53
I
2.5
85
At
2.1
In each pair of bonds, put a star () next to the more polar bond and use an arrow () to show
the direction of polarity in each separate bond.
2.
a)
C—O
and
C—N
c)
B—O
and
B—S
b)
P—Br
and
P—Cl
d)
B—F
and
B—I
For each of the bonds listed below, indicate () which atom is more negatively charged.
a)
C—N
b)
C—H
c)
C—Br
d)
S—O
It is somewhat artificial to classify bonds based on the differences in the electronegativities (X) of the two
atoms. However, we will use these ranges to do so:
Ionic
X > 1.7
(symbolized as A+ and Z-) – full charges
Polar Covalent
Nonpolar Covalent
3.
1.7  X  .5
X < .5
(symbolized as A+ and Z-) – partial charges
(no charges)
For each of the bonds listed below, classify each bond and indicate full, partial, or no charges.
a)
Na—Cl
e)
Mg—H
b)
C—O
f)
Cs—F
c)
Cu—O
g)
Cl—Cl
d)
C—H
h)
Al—Cl
C) Rules for Drawing Lewis Structures of Covalently Bonded Molecules (or Polyatomic Ions)
1) Sum the valence electrons from all atoms in the molecule or ion

For an ANION  add one electron to the total for each negative charge

For a CATION  subtract one electron from the total for each positive charge
2) Determine the central atom in the molecule and attach all other atoms to it

The central atom is the least electronegative element in the compound excluding hydrogen –
hydrogen can NEVER be central! (Note: typically, the central atom is the one written first in
the molecular formula)

Write the symbols for all the other atoms around the central atom. Connect the atoms to the
central with a single bond (a dash). Keep track of the electrons being used. Each single bond
made uses 2 electrons.

Chemical formulas are often written in the order in which the atoms are connected in the
molecule (ex: HCN  carbon is the central atom)
3) Complete the octets around all the outer atoms bonded to the central using nonbonding
electrons ( : )

Note: Hydrogen atoms may only ever have 2 electrons associated with them at any given time

Keep track of the electrons being used to complete the octets
4) Place any leftover electrons on the central atom, even if doing so results in more than an
octet of electrons around the atom.

Note: Expanded octets occur when more than 8 electrons surround an atom. This exception
to the octet rule is allowed for any atom in the 3rd row of the periodic table and after!

Keep track of the electrons – make sure the total number of valance electrons available were
used in the Lewis structure
5) If there are not enough electrons to give the central atom an octet, multiple bonds are
needed.
Single Bond (
):
Double Bond (
Triple Bond (
):
):

Remove one of the nonbonding pairs of electrons on one of the outer atoms and draw a
double bond (2nd dash) connecting the outer atom to the central. If need be, a triple bond (3rd
dash) may be formed by removing another nonbonding pair from the same outer atom.

Keep track of the electrons – make sure the total number of valance electrons available were
used in the Lewis structure
Drawing Lewis Structures – In Class Examples
Lewis Dot Structures:
1) Methane: CH4
4) Sulfite ion: SO32-
2) Hydrochloric Acid: HCl
5) Hydrocyanic Acid: HCN
3) Ammonia: NH3
6) Methanal (aka: formaldehyde): H2CO
Resonance Structures:
7) Nitrite ion: NO2 –
Exceptions to the Octet Rule:
1)
A central atom which does not have a
complete octet (FEWER than 8 electrons)
BF3
2) Expanded Octets  central atoms which
have MORE than 8 electrons (allowed only
for atoms in the 3rd period and after)
SBr6
Drawing Lewis Diagrams - Practice
# of valence e-‘s:________
1. SF4
3. PO43-
5. N2
# of valence e-‘s:________
# of valence e-‘s:________
7. CO32-
2. CH2F2
# of valence e-‘s:________
4. CBr4 # of valence e-‘s:________
6. CCl2O
# of valence e-‘s:________
# of valence e-‘s:________
8. ClO2-
10. H2S
12. NH4+
# of valence e-‘s:________
# of valence e-‘s:________
# of valence e-‘s:________
14. NO31-
9. OF2
# of valence e-‘s:________
11. C2H4 # of valence e-‘s:________
(attach carbons together – make symmetrical)
13. BCl3
# of valence e-‘s:________
# of valence e-‘s:________
15. IBr5
# of valence e-‘s:________
17. C2H2 # of valence e-‘s:________
(attach carbons together – make symmetrical)
19. PCl3 # of valence e-‘s:________
21. PF5
# of valence e-‘s:________
16. SO42-
# of valence e-‘s:________
18. CO # of valence e-‘s:________
20. SiO2
# of valence e-‘s:________
22. PH3 # of valence e-‘s:________
23. XeF4
# of valence e-‘s:________
25. CN- # of valence e-‘s:________
27. C2Br2 # of valence e-‘s:________
(attach carbons together – make symmetrical)
29. SeF4
# of valence e-‘s:________
24. SeH2
# of valence e-‘s:________
26. IF3 # of valence e-‘s:________
28. ClO-
# of valence e-‘s:________
30. AsF3 # of valence e-‘s:________