KEY
Practice Problems:
Bonding and Lewis Dot Structures
CHEM 1A
1. Calculate the difference in electronegativities for the following atoms, and state whether a bond
between them would be ionic, polar covalent, or non-polar covalent.
Difference in EN
a) Na & F
EN
F 4.0
Na 0.9
diff. 3.1
b) H & O
c) N & Br
EN
N 3.0
Br 2.8
diff. 0.2
d) Mg & O
Difference in EN
3.3
Type of Bond
ionic
EN
O 3.5
H 2.1
diff. 1.4
Ionic
1.7
Polar Covalent
0.4
0.0
Non-Polar
Covalent
polar covalent
non-polar covalent
EN
O 3.5
Mg 1.2
diff. 2.3
ionic
2. What two factors influence whether or not a molecule will be polar overall?
1) Polarity of individual bonds:
2) Molecular shape:
For the molecule to be polar overall, at least one bond must be polar
For the molecule to be polar overall, the polar bonds must be oriented so that their
dipoles do not cancel each other out
3. Give the valence level orbital notation and draw the Lewis dot structure for each of the following
neutral atoms and simple ions.
a) N
b) S
2s
2p
c) Ca2+
3s
3p
4s
4p
d) Se2–
4s
4p
4.
Note: Lewis structures do not need to reflect the
actual molecular shape. For example, CH4 could
be drawn with all four hydrogen atoms 90° from
Draw Lewis dot structures for the following molecules:
each other. The geometry is shown in this
 Using VSEPR theory, predict the molecular shape and bond angle key to assist you in your visualization of
the shapes and bond angles.
 State whether the molecule is polar or non-polar overall

If the molecule is polar overall, show the appropriate partial charges (+ and –) on polar bonds
a) CH4
b) CH2O
–
C
H
total
–
–
Val. e
4
(1)4
C
H
O
total
8
En
C 2.5
H 2.1
diff. 0.4
tetrahedral
109.5°
non-polar molecule
c) H2O
+
En
O 3.5
C 2.5
diff. 1.0
trigonal planar
OCH = 121°
~ 120°
HCH = 118°
polar molecule
–
Val. e–
N
5
H (1)3
total 8
+
En
O 3.5
H 2.1
diff. 1.4
12
En
C 2.5
H 2.1
diff. 0.4
d) NH3
–
Val. e–
O
6
H (1)2
total 8
Val. e
4
(1)2
6
+
bent (angular)
< 109.5° (104.5°)
polar molecule
+
+
trigonal pyramidal
< 109.5° (107.8°)
polar molecule
En
N 3.0
H 2.1
diff. 0.9
Note: You are only required to
know the approximate bond angles
(in bold). The exact bond angles are
provided only for curiosity’s sake.
5. Draw Lewis dot structures for the following polyatomic ions:



Show resonance structures where appropriate
Using VSEPR theory, predict the shape and bond angle
Include formal charges in at least one of the structures
formal charge = val. e– – lone e– – ½ bonding e–
a) NH4+
b) CO32–
0
Val. e–
N
5
H (1)4
charge
–1
total 8
tetrahedral
109.5°
c) NO2–
trigonal planar
120°
0
–1
0
–1
0
Val. e–
N
5
O (6)2
charge
1
total 18
0
–1
Val. e–
C
4
O (6)3
charge
2
total 24
0
+1
0
+
0
bent (angular)
< 120° (115.4°)
d) CN–
Val. e–
N
5
C
4
charge
1
total 10
linear
no bond angle (only 2 atoms)
–1
0
6. Three non-equivalent Lewis dot structures are shown below for the polyatomic ion thiocyanate
(SCN–).
a) Assign a formal charge to each of the atoms. The equation for calculating formal charge should
be shown, but work is not required for each.
formal charge = val. e– – lone e– – ½ bonding e–
–1
0
0
0
0
–1
+1
0
–2
b) Which one of the possible structures for SCN– shown would contribute the LEAST to the overall
structure? Explain why.
This structure contributes the least because it has larger formal charges on the atoms
c) Of the two remaining structures, which structure is likely to contribute MOST to the overall
structure of SCN–? Explain why.
This structure contributes the most because it has the negative formal charge on the
more electronegative atom