Formula
1
Lewis Diagram
2
Structural
Formula
3
Shape Diagram
4
Shape
Code(s) and
Shape
Name(s)
5
Polarity of
Bonds
(P, NP or
both)
6
Polarity of
Molecule
(P or NP)
7
Types of
Intermolecular
Forces
(LD, DD, HB)
1. NH3
2. CBr4
3. H2S
4. PCl3
5. SCl2
6. CH2Cl2
7. HBr
Chem 20 Chemical Bonding Workbook
1
Jill Agnew
Formula
1
Lewis Diagram
2
Structural
Formula
3
Shape Diagram
4
Shape
Code(s) and
Shape
Name(s)
5
Polarity of
Bonds
(P, NP or
both)
6
Polarity of
Molecule
(P or NP)
7
Types of
Intermolecular
Forces
(LD, DD, HB)
8. OCl2
9. NI3
10. O2
11. SBr2
12. HCl
13. H2Te
14. NF3
Chem 20 Chemical Bonding Workbook
2
Jill Agnew
Formula
1
Lewis Diagram
2
Structural
Formula
3
Shape Diagram
4
Shape
Code(s) and
Shape
Name(s)
5
Polarity of
Bonds
(P, NP or
both)
6
Polarity of
Molecule
(P or NP)
7
Types of
Intermolecular
Forces
(LD, DD, HB)
15. H2Se
16. CH3OH
17. CHCl3
18. Cl2
19. CS2
20. CH4
21. OF2
Chem 20 Chemical Bonding Workbook
3
Jill Agnew
Formula
1
Lewis Diagram
2
Structural
Formula
3
Shape Diagram
4
Shape
Code(s) and
Shape(s)
5
Polarity of
Bonds
(P, NP or
both)
6
Polarity of
Molecule
(P or NP)
7
Types of
Intermolecular
Forces
(LD, DD, HB)
22. C2H2
23. CO2
24. NCl3
25. C2H4
26. H2O
27. SiH4
28. CHCF
Chem 20 Chemical Bonding Workbook
4
Jill Agnew
Boiling Points and Number of Electrons for Groups 14-17 Hydrogen Compounds
Group
Compound
CH4
14
15
16
17
# of Electrons
Boiling Point
162C
SiH4
112C
GeH4
90C
SnH4
50C
NH3
33C
PH3
87C
AsH3
60C
SbH3
25C
H2O
100C
H2S
65C
H2Se
45C
H2Te
15C
HF
20C
HCl
85C
HBr
69C
HI
35C
1. Complete the “Number of Electrons” column.
2. Graph the data above on one set of axes. Label both axes and give it a title. The
manipulated variable (# of electrons) goes on the x-axis and the responding variable (boiling
point) goes on the y-axis. Join the points for each group separately with a straight line
between each point.
3.
Label each point on the graph with the compound formula.
4. The hydrogen compounds of Groups 15, 16 and 17 elements have consistently increasing
van der Waals forces (except for the first hydrogen compounds) with increasing number of
electrons. Explain why the first hydrogen compounds show a reversal in this trend.
5. Explain why CH4 does not show this same reversal in trend that is displayed by the first
hydrogen compounds in the other Groups.
6. Explain why the boiling points of the hydrogen compounds of the Group 14 elements are
consistently lower than the boiling points of the other hydrogen compounds.
Chem 20 Chemical Bonding Workbook
5
Jill Agnew
Molecular
Compound
Number
of
Electrons
Boiling
Point (C)
1. F2(g)
188
2. Cl2(g)
35
3. Br2(l)
59
4. I2(s)
84
5. ClF(g)
101
6. BrF(g)
20
7. BrCl(g)
5
8. ICl(g)
97
9. IBr(g)
116
10. CH4(g)
162
11. C2H6(g)
87
12. C3H8(g)
45
13. C4H10(g)
0.50
14. C5H12(l)
36
15. CF4(g)
129
16. CCl4(l)
77
17. CBr4(s)
189
18. CH3F(g)
78
19. CH3Cl(g)
24
20. CH3Br(g)
3.6
21. CH3I(l)
43
22. CH3OH(l)
65
23. C2H5F(g)
38
24. C2H5Cl(g)
13
25. C2H5Br(l)
38
26. C2H5I(l)
72
27. C2H5OH(l)
78
1. Complete the chart above.
Chem 20 Chemical Bonding Workbook
6
Jill Agnew
Types of Intermolecular Forces
London
DipoleHydrogen
Dispersion
Dipole
Bonding
2. Explain why there is a difference in boiling point between Br2(g) and ICl(g).
3. Explain why there is a difference in boiling point between BrF(g) and CH3F(g).
4. Why would boiling point increase if number of electrons increases?
5. Explain why methanol (CH3OH(l)) and ethanol (C2H5OH(l)) each have the lowest number of
electrons but the highest boiling point of their respective series.
6. Explain why there is a difference in boiling point between Cl2(g) and C4H10(g).
7. Explain why there is a difference in boiling point between BrCl(g) and C2H5Br(l).
Chem 20 Chemical Bonding Workbook
7
Jill Agnew
Chemical Bonding Review
1. Define the following terms:
a) valence level
b) electronegativity
c) ionic bond
d) metallic bond
e) covalent bond
f) lone pair
g) bonding electron
h) bonding capacity
i) crystal lattice
j) nonpolar covalent bond
k) polar covalent bond
l) polarity
m) intermolecular forces
n) intramolecular forces
o) dipole-dipole forces
p) hydrogen bonding
q) London Dispersion forces
2. Draw the electron dot diagrams for the following ionic compounds:
a) AlCl3
c) magnesium chloride
b) Na2S
d) calcium phosphide
3. How does electronegativity change down group and across period?
4. What two things affect electronegativity?
5. Compare the properties of ionic compound to the properties of molecular
compounds.
6. Draw the Lewis structure for the following:
e) PCl3
f) N2
g) CO2
h) SiH4
i) CH3OH
j) HCN
k) HNS
l) C2F4
m) HF
n) H2S
o) CI4
p) NF3
7. State the shape code and the shape for each of the molecules in question 6.
8. State the polarity of each molecule in question 6.
9. State the intermolecular attractions present for each of the molecules in question
6.
Chem 20 Chemical Bonding Workbook
8
Jill Agnew
10.Draw the bond dipole diagram for each of the following bonds and state
whether it is polar, nonpolar or ionic:
a) H – Br
e) H – N
b) S – S
f) O – Cl
c) Ca – I
g) F – F
d) C – O
h) C – H
11.Draw and label a “Scale of Forces” diagram.
12.Explain how conductivity relates to ionic compounds, metals and molecular
compounds.
Chem 20 Chemical Bonding Workbook
9
Jill Agnew