AP Chemistry
Unit 3 Review Guide
Name __________________________
Class _______ Date _______
Key Terms:
potential energy
exothermic
Coulomb’s law
covalent bond
Lewis structure
resonance
delocalized electrons
endothermic
ionic bonds
polar covalent
nonpolar covalent
formal charge
bond energy
molecular geometry
cation
anion
crystal lattice
metallic bond
electronegativity
VSEPR
sigma bonds
hybrid orbitals(sp, sp3, sp3d, etc.)
lattice energy
bond polarity
pi bonds
dipole
Equations:
𝐸=
1.
𝑄1 ∙𝑄2
𝑟
FC= valence e–– [nonbonding e- + ½ bonding e-]
Why do bonds form? What forces are involved in bonding?
∆𝐻𝑟𝑥𝑛 = Σ(𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛) + Σ(𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑)
10. What is lattice energy?
11. What two factors affect lattice energy? Write the equation.
2.
Compare and contrast ionic bonding and covalent bonding.
12. What is a crystal lattice?
3.
4.
What circumstance must exist for a bond to be nonpolar
covalent?
When atoms of a metal react with atoms of a nonmetal, what
type of electron configurations do the resulting ions attain?
13. For each of the following sets of elements, identify which
element would be expected to be most electronegative and
which would be expected to be least electronegative.
a) K, Sc, Ca
b) Br, F, At
c) C, O, N
5.
What are 3 basic properties of ionic compounds?
12. On the basis of the electronegativity values, indicate whether
each of the following bonds would be expected to be ionic,
nonpolar covalent, polar covalent, or metallic.
a) H-F
c) Cl-F
e) Cu-Zn
b) Na-F
d) Ca-F
f) Fe-Fe
6.
What are 3 basic properties of polar/nonpolar molecules?
13. On the basis of electronegativity values, indicate which is the
more polar bond in each of the following pairs.
a) H-N or H-P
c) H-P or H-S
b) H-O or H-S
d) H-S or H-I
7.
Explain the difference between single, double, and triple
bonds. Which is strongest? Which has the shortest bond
length?
14. In each of the following molecules, which end of the molecule
is negative relative to the other end? Use dipole arrows.
a) HCl
c) BrF
b) CO
d) HF
15. Consider the lattice energies of the following alkali metal
chlorides: Why does the magnitude of the lattice energy
decrease as we move down the group?
metal chloride
8.
Explain the difference between a covalent bond formed
between two atoms of the same element and a covalent
bond formed between atoms of two different elements?
LiCl
NaCl
KCl
CsCl
lattice energy
kJ/mol
-834
-788
-701
-657
16. Arrange these ionic compounds in order of increasing
magnitude of lattice energy: CaO, KBr, KCl, SrO.
9.
What is the “glue” or force that holds compounds (ionic)
together. How is that different from the “glue” that holds
molecules (covalent) together?
1
17. Use the average bond energy table in your notes to calculate
the overall enthalpy change (∆𝐻𝑟𝑥𝑛 ) of the following reaction:
(Lewis structures will help)
CH4 + Cl2 → CH3Cl + HCl
20. Assign formal charges and determine which structure is most
stable.
21. Give the name or write the formula for the following:
a. N2O
____________________
18. Examine the graph and determine the bond energy and bond
length for the following two bonded atoms:
b.
Cr(NO2)2
____________________
c.
V2O5
____________________
d.
(NH4)2Cr2O7
____________________
e.
phosphorous trioxide ____________________
f.
copper (III) silicate
____________________
g.
beryllium nitride
____________________
19. Draw the Lewis structures and determine the formal charge
for each atom in N2O. Circle the most stable resonance
structure.
22. Complete the chart.
Lewis Structure
Molecular
geometry
Bond Angles
Hybridization (sp,
sp3, etc.)
Molecular
Polarity
SCN−
O3
BCl3
IF5
SO2
2
23. Explain why CO2 and CCl4 are both nonpolar even though
they contain polar bonds?
24. The structure of caffeine is shown. How many pi and sigma
bonds are present.
25. Explain why BF3 is trigonal planar, yet NH3 is trigonal
pyramid.
26. Why does electronegativity increase as you move across a
period?
27. Why is the bond angle for H2O 109°, yet for SO2
slightly < than 120°?
28. Formaldehyde, CH2O is shown. Explain why the bond angles
are different.
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