AP Chemistry
Unit 3 Practice Problems
1.
Name _________________________________
Date ____________ Class _____________
Consider the following potential energy diagram of two
different bonds formed between different atoms.
5.
Draw the Lewis structure and determine whether each of the
following molecules are polar or nonpolar.
H2O
BF3
a.
What happens to the potential energy as the atoms get
closer to each other and form a bond? Explain.
H2S
b.
Does this difference in potential energy cause a release of
energy or absorption of energy? Explain.
CCl4
c.
What happens to the potential energy if they get close?
Explain.
C2H4
d.
Which bond will release more energy when it forms?
Explain.
e.
Which bond has a shorter bond length? Explain.
CO2
2.
Use the average bond energy table in your notes to
calculate the amount of energy required to completely
separate an acetylene molecule, C2H2, into individual atoms.
N2
NH3
3.
Use the average bond energy table in your notes to
calculate the overall enthalpy change (∆𝐻𝑟𝑥𝑛 ) of the
following reaction: (Lewis structures will help)
C2H4 + H2 → C2H6
6.
4.
Determine the bond type:
a. C and S
b. N and S
c. Na and F
d. N and O
e. O and O
f. H and C
Consider the main group elements (1-2, 13-18).
a. Record the number of valence electrons.
b. Draw the Lewis dot structure for element "X"
c. Record the ionic charge when forming ionic bond
d. Record the total number of electrons surrounding the
atom when forming covalent bond(s)
1
2
13
14
15
16
17
18
a
b
X
X
X
X
X
X
X
X
c
d
1
7.
Illustrated below are four ions—A+, B+, C- and D-—showing
their relative ionic radii.
A+
B+
C–
D–
Indicate whether the following correlate directly or inversely.
Direct Inverse
Electronegativity difference & dipole moment
Electronegativity difference & bond strength
Dipole moment & bond strength
Bond length & bond strength
What combinations of ions are impossible?
What would have the greatest lattice energy?
What would have the least lattice energy?
8.
The table lists the ionic radius (x 10-10 m) of common ions.
Li+ (0.68) Be2+ (0.31)
O2- (1.40)
F- (1.33)
+
2+
3+
2Na (0.97) Mg (0.66) Al (0.51) S (1.84) Cl- (1.81)
a. Use the above information to estimate the relative
lattice energy for each ionic bond. E  Q1Q2/d
(Q1 and Q2 = ionic charge and d  (rcation + ranion)
Ionic Bond
Relative Lattice Energy
11. Consider the following data for C-C bonds.
C-C bond
Single
Double
Triple
Bond Strength (kJ/mol)
348
614
839
How does bond strength correlate with the number of
shared electrons?
12. Complete the chart with the formula or name of the binary
molecule.
Formula
Name
LiF
N2O5
MgO
Fe(NO3)2
carbon tetrachloride
NaCl
CO2
Al2S3
b.
copper (II) sulfate
Ionic compounds melt when the temperature is high
enough to break the ionic bond. Rank the above
compounds in order of lowest to highest melting point.
nitrogen monoxide
OF2
c.
9.
What is the relative lattice energy for the strongest ionic
bond formed from the ions listed in the table?
Use the electronegativity values to answer the questions.
H 2.1
Li 1.0 Be 1.5
B 2.0
C 2.5
N 3.0
O 3.5 F 4.0
Na 0.9 Mg 1.2 Al 1.5 Si 1.8
P 2.1
S 2.5 Cl 3.2
K 0.8 Ca 1.0 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8
Rb 0.8 Sr 1.0
In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5
a. What is the range of electronegativities?
metals
metalloids
nonmetals
b.
aluminum phosphate
13. Record the number of covalent bonds typically formed by the
main group elements.
1
2
13
14
15
16
17
18
14. In the Lewis structure shown below, A, D, E, Q, X and Z
represent non-metal elements in the first two rows of the
periodic table. Identify the elements.
Rank the following bonds from most polar (1) to least
polar (6). Place + next to the atom with the lower
electronegativity.
N–S
O–S
F–S
P–S
S–S
10. Consider the following data for hydrogen halides.
Bond
HElectronegativity
Dipole
Strength
Halide
difference
moment
(kJ/mol)
HF
1.9
1.82
436
HCl
0.9
1.08
431
HBr
0.7
0.82
366
0.4
0.44
299
HI
Cl–S
A
D
E
Q
X
Z
15. There are three ways to draw Lewis structures for NCO–
a. Calculate the formal charge for each version
Structure [:::N–CO:]–
[::N=C=O::]–
[:NC–O:::]–
Formal
Charge
Bond
Length
(Å)
0.92
1.27
1.41
1.61
b.
Which is the preferred structure? Give two reasons.
2
16. Draw Lewis structures for the following molecules where the
first atom listed is the central atom, unless indicated.
POCl3 (lowest formal charge) CNO– (lowest formal charge)
SCN– (C is central atom)
18. Label the hybridization for each carbon atom.
IF4–
19. Draw the resonance structures for the following molecules.
Molecule
Resonance Structures
SF6
CO32-
BrF3
SO2
SO3
N2O (lowest formal charge)
OCN–
20. Given the following species: CO32-, CO2, CO.
a. Draw the Lewis structure for each.
IF5
NO2
SF4
SO2Cl2 (lowest formal charge)
b.
Which one has the shortest C—O bond length? Explain.
17. Determine the number of  bonds, the number of  bonds,
number of lone electron pairs, hybridization around the
central atom.
Lone


Molecule
Hybridization
Pairs
Bonds Bonds
CH4
NH3
CO2
CH2O
POCl3
CNO–
SCN–
BCl3
SF6
BrF3
HF
SO3
N2 O
IF5
SO2Cl2
XeF4
3
c.
4