Chemistry Practice
Name ___________________
Chapter 8: Chemical Bonding
Concept Development Exercises
1.
Distinguish between ionic, covalent, and metallic bonding. Give examples of
substances, too, that employ these kinds of bonds.
Type of Bond
Definition
Examples
(substances)
ionic bonding
covalent bonding
metallic bonding
2.
Draw Lewis symbols for atoms of nitrogen, oxygen, fluorine, and neon.
3.
State the octet rule:
4.
Metals and nonmetals combine to form ionic compounds. Compare the properties of
these elements that make them well-suited for producing ionic compounds.
Element
Type
metal
nonmetal
Atomic Size
Number of
valence e-s
Ionization
Energy
Electron
Affinity
Lose or
Gain e-s
5.
A. Define lattice energy:
B. What factors determine the magnitude of an ionic compound’s lattice energy?
C. What does the symbolism, H, mean?
True or False
6. For a given arrangement of ions, the lattice energy increases as the charges on the
ions increase and as their radii decrease.
7.
Arrange the following compounds in order of increasing lattice energy: NaF, CsI,
and CaO.
8.
A. Write the electron configurations for Ca and Ca+2.
B. Even though lattice energy increases with increasing ionic charge, we never find
ionic compounds that contain Ca+3 ions . Explain.
True or False
9. Transition-metal ions generally do not form ions that a noble gas electron
configuration.
10. Specifically, from which orbitals are electrons lost when an iron atom, Fe, becomes
an iron (III) ion, Fe+3? (Begin by writing their electron configurations.)
11. Write empirical formulas for compounds that contain the following ions:
A. NH4+1; CO3-2
B. Sr+2; Cl-1
C. Al+3; SO4-2
D. Fe+2; PO4-3
12. The reaction between sodium metal and chlorine gas is exothermic (H = -411 kJ).
Follow the Born-Haber cycle for NaCl below (and p. 308, text) and verify the
enthalpy change for the reaction: Na(s) + ½ Cl2 (g)  NaCl(s)
1. vaporization of Na(s) to Na(g):
2. atomization of Cl2 (g) to Cl(g):
3. ionization of Na(g) to Na+1(g):
4. ionization of Cl(g) to Cl-1(g):
5. Na+1(g) and Cl-1(g) to NaCl(s) lattice:
13. Identify several properties of ionic substances:
True or False
14. The vast majority of chemical substances do not have the characteristics of ionic
substances.
15. Specifically, what holds the hydrogen atoms together in an H2 molecule?
16. Draw Lewis models for the following covalent compounds: HF, H2O, NH3, and
CH4.
17. Draw Lewis models for CO2 and N2.
True or False
18. As a general rule, the distance between bonded atoms decreases as the number of
shared electrons increases.
19. Distinguish between a polar covalent bond and a nonpolar covalent bond.
20. Define electronegativity:
21. Identify the electronegativity trends in the periodic table (from left to right along a
period and downward in a group).
True or False
22. The greater the difference in electronegativity between two bonded atoms, the more
polar their bond.
23. Define dipole:
24. If two equal and opposite charges, Q+ and Q- are separated by a distance r, the
magnitude of the dipole moment is the product of Q and r:  = Qr. Dipole
moments are measured in debyes (D), a unit that equals 3.34 x 10-30 C-m. Work
through Sample Exercise 8.5 on pp. 315-316 in the text.
Drawing Lewis Structures
Lewis structures can help us understand the bonding in many compounds and are
frequently used when discussing the properties of molecules. Drawing Lewis structures is
an important skill that you should practice.
25. Draw Lewis structures for the following:
A. PCl3
B. HCN
C. NO+1
26. Draw Lewis structures for the following:
A. BrO3-1
B. ClO2-1
C. PO4-3
27. A. Define formal charge:
B. How do we calculate formal charge?
28. Determine the formal charges associated with each atom associated with the Lewis
structures from questions 26 and 27 above.
PCl3
HCN
NO+1
BrO3-1
ClO2-1
PO4-3
True or False
29. The sum of the formal charges will always equal the overall charge on an ion.
30. The concept of formal charge can help us choose between alternative Lewis
structures. For example, carbon dioxide could feature two double bonds in the
molecule or a single and a triple covalent bond between the carbon atom and the
oxygen atoms (see p. 320, bottom).
A. What guidelines do we follow to choose the most correct Lewis structure?
B. The cyanate ion, NCO-1, has three possible Lewis structures. Draw these three
Lewis structures and assign formal charges to the atoms in each structure.
C. Which of the three structures is preferred? Explain.
31. A. What are resonance structures?
B. The O-O bonds in ozone, O3, are often described as “one-and-a-half” bonds. Is
this description consistent with the idea of resonance?
C. Draw resonance structures for benzene, C6H6.
Exceptions to the Octet Rule
The octet rule fails in many situations involving covalent bonding, and are of three main
types:
1. Molecules and polyatomic ions containing an odd number of electrons;
2. Molecules and polyatomic ions in which an atom has fewer than an octet of
valence electrons;
3. Molecules and polyatomic ions in which an atom has more than an octet of
valence electrons.
32. Draw Lewis structures for nitric oxide, NO, boron trifluoride, BF3, phosphorus
pentachloride, PCl5, and the tetrachloroiodide ion, ICl4-1.
Bond Enthalpies
33. Calculate the enthalpy change for the following gas-phase reaction:
CH4 (g) + Cl2 (g)  CHCl3 (g) + HCl (g)
34. Nitroglycerin, C3H5N3O9 (l), is an explosive compound that decomposes into
nitrogen, carbon dioxide, water, and oxygen gases. Why is this reaction so
energetic?
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