Download CHAPTER 6 REVIEW Teacher Notes and Answers

CHAPTER 6 REVIEW
Teacher Notes and Answers
1.In ionic bonding, large numbers of oppositely
charged ions join because of mutual electrical
attraction. In covalent bonding, atoms join by
sharing electron pairs. In metallic bonding,
atoms join through an attraction to a sea of
valence electrons.
2.The electrons in a covalent bond occupy
overlapping orbitals; each electron is free to
occupy either of the orbitals, but both are more
likely to be in the space between the nuclei of
the bonded atoms.
3.A multiple bond is needed when there are not
enough valence electrons to complete octets by
adding unshared pairs.
4.Ionic compounds have higher melting points and
boiling points than molecular compounds do,
and they do not vaporize at room temperature.
5.Properties include hardness, brittleness, and
electrical conductivity in the molten state.
6.Metals are better heat conductors than ionic or
molecular compounds. Metals are more easily
deformed and are better electrical conductors in
the solid state. Metals are also shiny.
7.According to VSEPR theory, the shapes of
molecules are classified based on the number
of bonding electron pairs and lone pairs that
surround a molecule’s central atom.
8.Intermolecular forces are the forces of attraction
between molecules.
9.H and I: 2.4 − 2.0 = 0.4, polar-covalent, I;
S and O: 3.5 − 2.5 = 1.0, polar-covalent, O; K
and Br: 2.8 − 0.8 = 2.0, ionic, Br;
Si and Cl: 3.0 − 1.8 = 1.2, polar-covalent, Cl; K
and Cl: 3.0 − 0.8 = 2.2, ionic, Cl;
Se and S: 2.5 − 2.4 = 0.1,
nonpolar-covalent, S;
C and H: 2.4 − 2.0 = 0.4, polar-covalent, C
10.K and Cl, K and Br, Si and Cl, S and O, H and I
and C and H, Se and S
11.Bonding is stronger between the ions in sodium
chloride because its lattice energy is greater
(more negative). Greater lattice energy indicates
stronger ionic bonding.
12a. Cl
12b. Ca
12c. C
12d. P
F
FCF
13a. F
13b. H Se H
H
H C Cl
13c. H
13d. N N
14a.trigonal-pyramidal
14b.bent or angular
14c.bent or angular
15a.toward F
15b.toward Cl
15c.toward I
16a.nonpolar
16b.polar
16c.polar
16d.nonpolar
16e.polar
16f.polar
17a.polar
17b.nonpolar
17c.nonpolar
17d.polar
17e.nonpolar
18a. Cl S Cl bent or angular
I PI
18b. I
trigonal-pyramidal
18c. Cl O Cl bent or angular
HNH
18d. Cl trigonal-pyramidal
Cl
Cl Si Cl
Br
18e.
tetrahedral
18f. O N Cl bent or angular
-
18g.
ON O
O
O
O S O
O
18h.
trigonal-planar
2-
tetrahedral
CHEMICAL BONDING
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CHAPTER 6 REVIEW
1. Identify and define the three major types of chemical bonding.
2. Describe the general location of the electrons in a covalent bond.
3. In writing Lewis structures, how is the need for multiple bonds
generally determined?
4. In general, how do ionic and molecular compounds compare in terms of
melting points, boiling points, and ease of vaporization?
5. List three general physical properties of ionic compounds.
6. How do the properties of metals differ from those of both ionic and
molecular compounds?
7. How is VSEPR theory used to classify molecules?
8. What are intermolecular forces?
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CHAPTER 6
9. Complete the table with respect to bonds formed between the pairs of atoms.
Bonded atoms
Electronegativity difference
Bond type
More-negative atom
H and I
S and O
K and Br
Si and Cl
K and Cl
Se and S
C and H
10. List the bonding pairs from Question 9 in order of increasing
covalent character.
11. The lattice energy of sodium chloride, NaCl, is –787.5 kJ/mol. The lattice
energy of potassium chloride, KCl, is –715 kJ/mol. In which compound is the
bonding between ions stronger? Why?
12. Use electron-dot notation to illustrate the number of valence electrons
present in one atom of each of the following elements.
a. chlorine, Cl
c. carbon, C
b. calcium, Ca
d. phosphorus, P
13. Draw Lewis structures for each of the following molecules.
a. C​F4​ ​ c. CCl​H3​ ​ b. ​H2​ ​Se d. ​N​2​ CHEMICAL BONDING
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14 According to the VSEPR theory, what molecular geometries are associated
with the following types of molecules?
a. A​B3​ ​E
b. A​B2​ ​​E​2​
c. A​B2​ ​E
15. For each of the following polar molecules, in which direction is the dipole
oriented? (Which end would the dipole arrow point toward?)
a. H–F
b. H–Cl
c. H–I
16. Determine whether each of the following bonds would be polar or nonpolar.
a. H–H
c. H–F
e. H–Cl
b. H–O
d. Br–Br
f. H–N
17. On the basis of individual bond polarity and orientation, determine whether
each of the following molecules would be polar or nonpolar.
a. ​H​2​O
c. C​F4​ ​
b. ​I2​ ​
d. N​H3​ ​
e. C​O2​ ​
18. Draw Lewis structures for each of the following molecules or ions, and then
use VSEPR theory to determine the geometry of each.
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a. SC​l2​ ​ e. SiC​l3​ ​Br b. P​I3​ ​ f. ONCl c. C​l2​ ​O g. NO​ 
   3 ​ d. N​H2​ ​Cl  
h. SO​24  ​    CHAPTER 6