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IONS
7.1
Section Review
Objectives
• Determine the number of valence electrons in an atom of a representative
element
• Explain the octet rule
• Describe how cations form
• Explain how anions form
Vocabulary
• valence electrons
• electron dot structures
• octet rule
• halide ions
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
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Elements within the same group of the periodic table behave
1
similarly because they have the same number of
2
1.
2.
. The
number of a representative element indicates how many
3.
valence electrons that element has. Diagrams that show valence
4.
3
5.
electrons as dots are called
. Gilbert Lewis’s
4
states
that in forming compounds, atoms tend to achieve the electron
6.
configuration of a noble gas.
7.
The transfer of valence electrons produces positively charged
5
ions, or
, and negatively charged ions called
cations of Group 1A elements always have a charge of
6
7
. The
9.
.
10.
8
are produced when atoms of the elements in Group 7A
9
an electron. For transition metals, the
10
8.
of cations
may vary.
Chapter 7 Ionic and Metallic Bonding
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Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 11. The chlorine atom gains seven electrons when it becomes an ion.
________ 12. The chemical properties of an element are largely determined by the
number of valence electrons the element has.
________ 13. Atoms acquire the stable electron structure of a noble gas by losing
electrons.
________ 14. An atom of an element in Group 1A has seven valence electrons.
________ 15. Among the Group 1A and 2A elements, the group number of each
element is equal to the number of valence electrons in an atom of
that element.
________ 16. Sulfur and magnesium both have two valence electrons.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
a. ions that are produced when halogens gain electrons
________ 18. valence electron
b. a depiction of valence electrons around the symbol
of an element
________ 19. octet rule
c. has the electron configuration of argon
________ 20. cations
d. an electron in the highest occupied energy level of
an element’s atom
________ 21. anions
e. Atoms in compounds tend to have the electron
configuration of a noble gas.
________ 22. halide ions
f. atoms or groups of atoms with a negative charge
________ 23. chloride ion
g. atoms or groups of atoms with a positive charge
Part D Questions and Problems
Answer the following in the space provided.
24. Write the electron dot structures for the following atoms.
a. silicon
b. rubidium
c. barium
156
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________ 17. electron dot structure
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25. State the number of electrons lost or gained in forming each of these ions.
Name the ions and tell whether it is an anion or a cation.
a. Mg2!
c. Br"
b. Ca2!
d. Ag!
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26. Describe the formation of an ion from a metal and a nonmetal in terms of the octet rule.
Chapter 7 Ionic and Metallic Bonding
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7.2
Date ___________________
Class __________________
IONIC BONDS AND IONIC COMPOUNDS
Section Review
Objectives
• Explain the electrical charge of an ionic compound
• Describe three properties of ionic compounds
Vocabulary
• ionic compounds
• ionic bonds
• chemical formula
• formula unit
• coordination number
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1
Anions and cations attract one another by means of
The forces of attraction that hold
3
1.
charged ions together in
2.
. Although they are composed
3.
of ions, ionic compounds are electrically
4
. The lowest whole- 4.
5
number ratio of ions in an ionic compound is called a
6
Nearly all ionic compounds are solid
at room
temperature. Ionic compounds in general have very
8
melting temperatures. This is because the
9
7
attractive
.
5.
6.
7.
8.
structure.
9.
Ionic compounds conduct an electric current when in the
10.
forces between the ions result in a very
10
state or dissolved in water.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 11. During the formation of the compound NaCl, one electron is
transferred from a sodium atom to a chlorine atom.
158
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ionic compounds are called
2
.
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________ 12. The coordination number of an ion is the number of ions of positive
charge that surround the ion in a crystal.
________ 13. The coordination number of the ion Na! in NaCl is 6.
________ 14. In forming an ionic compound, an atom of an element gains electrons.
________ 15. Ionic compounds cannot conduct electricity if they are dissolved in water.
Part C Matching
Match each description in Column B to the correct term in Column A.
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Column A
Column B
________ 16. ionic compounds
a. the number of ions of opposite charge surrounding
each ion in a crystal
________ 17. ionic bonds
b. compounds composed of cations and anions
________ 18. chemical formula
c. shows the kinds and numbers of atoms in the
smallest representative unit of a substance
________ 19. formula unit
d. lowest whole-number ratio of ions in an ionic
compound
________ 20. coordination number
e. the electrostatic forces of attraction binding
oppositely charged ions together
Part D Questions and Problems
Answer the following in the space provided.
21. List the characteristics of an ionic bond.
22. Explain the electrical conductivity of melted and of aqueous solutions of ionic
compounds using the characteristics of ionic compounds.
Chapter 7 Ionic and Metallic Bonding
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BONDING IN METALS
7.3
Section Review
Objectives
• Model the valence electrons of metal ions
• Describe the arrangement of atoms in a metal
• Explain the importance of alloys
Vocabulary
• metallic bonds
• alloys
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
Metals consist of closely packed
2
by a sea of
1
that are surrounded
. This arrangement constitutes the
3
4
conductivity of metals and helps explain why
5
metals are
7
packed in a
and
6
cubic, a
2.
3.
4.
. Metal atoms are commonly
5.
8
6.
cubic, or a
9
arrangement. When two or more elements, at least one of which
7.
is a metal, are mixed together, the resulting mixture is called
8.
an
10
.
9.
10.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 11. In a body-centered cubic structure, each atom has 12 neighbors.
________ 12. Metallic objects are formed from pure metals.
160
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bond. The electron mobility accounts for the excellent
1.
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________ 13. Metals that are good conductors of electricity are said to be ductile.
________ 14. Drifting valence electrons insulate cations from one another and
contribute to the malleability of a metal.
________ 15. Metals are good conductors of electricity because electrons can flow
freely in them.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
________ 16. ductile
a. an alloy whose component atoms are different sizes
________ 17. metallic bonds
b. a mixture of two or more elements, at least one of which
is a metal
________ 18. alloy
c. can be hammered or forced into shapes
________ 19. malleable
d. can be drawn into wires
________ 20. interstitial alloy
e. the attraction of valence electrons for positive metal ions
Part D Questions and Problems
Answer the following in the space provided.
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21. Explain the physical properties of metals, using the theory of metallic bonding.
22. Explain why the properties of alloys are generally superior to their constituent
components.
Chapter 7 Ionic and Metallic Bonding
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7
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IONIC AND METALLIC BONDING
Practice Problems
In your notebook, answer the following.
SECTION 7.1 IONS
1. For each element below, state (i) the number of valence electrons in the atom,
(ii) the electron dot structure, and (iii) the chemical symbol(s) for the most
stable ion.
a. Ba
b. I
c. K
2. How many valence electrons does each of the following atoms have?
a. gallium
b. fluorine
c. selenium
3. Write the electron configuration for each of the following atoms and ions.
e. O2"
a. Ca
c. Na!
b. chlorine atom
d. phosphide ion
4. What is the relationship between the group number of the representative
elements and the number of valence electrons?
5. How many electrons will each element gain or lose in forming an ion? State
whether the resulting ion is a cation or an anion.
c. tellurium
e. bromine
b. aluminum
d. rubidium
f. phosphorus
6. Give the name and symbol of the ion formed when
a. a chlorine atom gains one electron.
b. a potassium atom loses one electron.
c. an oxygen atom gains two electrons.
d. a barium atom loses two electrons.
7. How many electrons are lost or gained in forming each of the following ions?
a. Mg2!
b. Br"
c. Ag!
8. Classify each of the following as a cation or an anion.
162
a. Na!
c. I"
e. Ca2!
b. Cu2!
d. O2"
f. Cs!
Core Teaching Resources
d. Fe3!
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a. strontium
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SECTION 7.2 IONIC BONDS AND IONIC COMPOUNDS
1. Use electron dot structures to predict the formula of the ionic compounds formed
when the following elements combine.
a. sodium and bromine
d. aluminum and oxygen
b. sodium and sulfur
e. barium and chlorine
c. calcium and iodine
2. Which of these combinations of elements are most likely to react to form ionic
compounds?
a. sodium and magnesium
c. potassium and iodine
b. barium and sulfur
d. oxygen and argon
3. What is the meaning of coordination number?
4. How is the coordination number determined?
SECTION 7.3 BONDING IN METALS
1. What is a metallic bond?
2. How is the electrical conductivity of a metal explained by metallic bonds?
3. Are metals crystalline? Explain.
4. Give three possible crystalline arrangements of metals. Describe each.
5. What is an alloy?
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6. Name the principal elements present in each of the following alloys.
a. brass
d. sterling silver
b. bronze
e. cast iron
c. stainless steel
f. spring steel
Chapter 7 Ionic and Metallic Bonding
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8.1
Date ___________________
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MOLECULAR COMPOUNDS
Section Review
Objectives
• Distinguish molecular compounds from ionic compounds
• Identify the information a molecular formula provides
Vocabulary
• covalent bond
• molecule
• diatomic molecule
• molecular compound
• molecular formula
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
Every substance is either an element or a(n)
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A compound is either
2
.
or ionic in nature. Most molecular
3
4
consisting of two atoms are
ionic compounds tend to have
molecules. The chemical
4.
5
. Molecular
5.
melting and boiling points, while
6.
7
7.
melting and boiling points.
A molecular formula shows how many
8
of each
element a molecule contains, but it does not indicate the
9
2.
3.
formula of a molecular compound is a
6
1.
. Molecules
compounds are composed of two or more
compounds tend to have
1
8.
9.
of the molecule.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 10. A diatomic molecule contains two or three atoms.
________ 11. Molecular compounds have relatively high boiling points.
Chapter 8 Covalent Bonding
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________ 12. The molecular structure of carbon dioxide is one carbon atom with
two oxygen atoms on opposite sides of it.
________ 13. Covalent bonds exist when combining atoms give up or accept electrons.
________ 14. A molecule contains two atoms.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
________ 15. molecule
a. compound composed of molecules
________ 16. molecular compound
b. a molecule consisting of two atoms
________ 17. covalent bond
c. shows the kinds and numbers present in a molecule of
a compound
________ 18. diatomic molecule
d. joins atoms held together by sharing electrons
________ 19. molecular formula
e. an electrically neutral group of atoms joined together
by covalent bonds
Part D Questions and Problems
Answer the following in the space provided.
21. Identify the number and kinds of atoms present in a molecule of each
compound.
a. butane (C4H10)
b. fluorobenzene (C6H5F)
22. Classify each particle as an atom or a molecule.
a. CH4
d. He
b. Ne
e. CO2
c. O2
182
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20. A compound has a boiling point of 40°C. Is this compound most likely an ionic
or a molecular compound?
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THE NATURE OF COVALENT BONDING
8.2
Section Review
Objectives
•
•
•
•
•
•
•
State a rule that usually tells how many electrons are shared to form a covalent bond
Describe how electron dot formulas are used
Predict when two atoms are likely to be joined by a double or a triple covalent bond
Distinguish between a single covalent bond and other covalent bonds
Describe how the strength of a covalent bond is related to its bond dissociation energy
Describe how resonance structures explain bonding
Identify some exceptions to the octet rule
Vocabulary
• single covalent bond
• structural formulas
• unshared pairs
• double covalent bonds
• triple covalent bonds
• coordinate covalent bond
• polyatomic ion
• bond dissociation energy
• resonance structures
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Part A Completion
Use this completion exercise to check your understanding of the concepts and terms that are
introduced in this section. Each blank can be completed with a term, short phrase, or number.
1
When atoms share electrons to gain the
of a noble gas, the bonds formed are
4
2
3
.A
configuration
1.
pair of
2.
covalent bond. Pairs of
3.
valence electrons that are not shared between atoms are called
4.
valence electrons constitutes a
5
5.
. Sometimes two or three pairs of electrons may be shared
6
to give
6.
covalent bonds. In some cases, only one of the
7.
atoms in a bond provides the pair of bonding electrons; this is a
7
.
8
8.
is required to break covalent bonds between
9.
atoms. The total energy required to break the bond between two
covalently bonded atoms is known as the
9
10.
.
When it is possible to write two or more valid electron dot
formulas for a molecule or ion, each formula is referred to as a
10
.
Chapter 8 Covalent Bonding
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Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 11. The modern interpretation of resonance is that electron pairs rapidly flip
back and forth between the various electron dot structures.
________ 12. The compound NH3 contains two double covalent bonds.
________ 13. The chemical formulas of molecular compounds show the number and
type of atoms in each molecule.
________ 14. A molecule of bromine has six unshared pairs of electrons.
________ 15. Carbon forms four single covalent bonds with other atoms.
________ 16. A bond in which one atom contributes both bonding electrons is called a
polyatomic covalent bond.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
a. a chemical formula that shows the arrangement of
atoms in a molecule or a polyatomic ion
________ 18. structural formula
b. the amount of energy required to break a covalent bond
between atoms
________ 19. bond dissociation
energy
c. a tightly bound group of atoms that has a positive or
negative charge and behaves as a unit
________ 20. polyatomic ion
d. a covalent bond in which one atom contributes both
bonding electrons
________ 21. coordinate covalent
bond
e. a chemical bond in which only one pair of electrons is
shared by two bonded atoms
Part D Questions and Problems
Answer the following in the space provided.
22. Draw electron dot structures for each of the following compounds
a. Br2
b. HCN
c. NH4!
184
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________ 17. single covalent bond
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Date ___________________
Class __________________
BONDING THEORIES
Section Review
Objectives
• Identify the difference between atomic and molecular orbits
• Describe how VSEPR theory helps predict the shapes of molecules
• Identify the ways in which orbital hybridization is useful in describing molecules
Vocabulary
• molecular orbitals
• bonding orbital
• sigma bond
• pi bond
• tetrahedral angle
• VSEPR theory
• hybridization
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
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The quantum mechanical model of bonding assumes that
atomic orbitals overlap to produce
1
. A molecular orbit that
can be occupied by two electrons of a covalent bond is called a
2
3
1.
2.
3.
than that of the atomic orbitals
4.
from which it formed. When two atomic orbitals combine to form
5.
a molecular orbital that is symmetrical around the axis connecting
6.
, whose energy is
4
two atomic nuclei, a
bond is formed. When atomic
orbitals overlap side by side, they produce
Electron dot structures fail to reflect the
of molecules.
7
5
bonds.
6
7.
8.
shapes
states that because electron pairs repel,
molecular shape adjusts so the valence-electron pairs are as
far apart as possible. Another way to describe molecules that
provides information about both molecular bonding and
molecular shape is
8
.
Chapter 8 Covalent Bonding
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Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 9. Unshared pairs of electrons affect the shape of molecules.
________ 10. Molecular orbitals involve pi bonding.
________ 11. A bonding orbital is a molecular orbital whose energy is higher than that
of the atomic orbitals from which it is formed.
________ 12. With hybridization, several atomic orbitals overlap to form the same total
number of equivalent hybrid orbitals.
________ 13. Sigma and pi bonds are found in the same molecule.
________ 14. The methane molecule has four orbitals with tetrahedral angles of 109.5°.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
a. states that because electron pairs repel, molecules adjust their
shapes so that valence-electron pairs are as far apart as possible
________ 16. pi bond
b. a process in which several atomic orbitals overlap to form the
same number of equivalent hybrid orbitals
________ 17. VSEPR theory
c. a term used to describe the shape of certain molecules such as
CO2
________ 18. hybridization
d. a bond formed when two atomic orbitals combine to form a
molecular orbital that is symmetrical along the axis connecting
the two atomic nuclei
________ 19. linear molecule
e. a bond in which the bonding electrons are most likely to be
found in the sausage-shaped regions above and below the
nuclei of the bonded atoms
Part D Questions and Problems
Answer the following in the space provided.
20. Indicate the hybrid orbitals used by each carbon atom in the following
compound.
2
2
H3C 2 C 3 C 2 C 4 C 2 CH3
H
186
Core Teaching Resources
H
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________ 15. sigma bond
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POLAR BONDS AND MOLECULES
8.4
Section Review
Objectives
• Describe how electronegativity values determine the charge distribution in a polar
bond
• Describe what happens to polar molecules when placed between oppositely
charged metal plates
• Distinguish intermolecular attractions from ionic bonds and from covalent bonds
• Identify the reason network solids have high melting points or decompose
without melting
Vocabulary
•
•
•
•
• dipole
• van der Waals forces
• dipole interactions
nonpolar covalent bond
polar covalent bond
polar bond
• dispersion forces
• hydrogen bonds
• network solids
polar molecule
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Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1.
When like atoms are joined by a covalent bond, the bonding
electrons are shared
1
, and the bond is
2
2.
. When the
atoms in a bond are not the same, the bonding electrons are shared
3
4
, and the bond is
. The degree of polarity of a bond
. The attractions between opposite poles of polar molecules
are called
is the
8
7
6
. Another strong intermolecular attractive force
, in which a hydrogen covalently bonded to a very
atom, such as
9
, is also weakly bonded to an
4.
5.
between any two atoms is determined by consulting a table of
5
3.
6.
7.
8.
9.
unshared electron pair of another electronegative atom.
Chapter 8 Covalent Bonding
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Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 10. In a polar covalent bond, the more electronegative atom has a slight
positive charge.
________ 11. In general, the electronegativity values of nonmetallic elements are
greater than the electronegativity values of metallic elements.
________ 12. A molecule with polar bonds is dipolar.
________ 13. Covalent compounds are network solids.
________ 14. If the electronegativity difference between two atoms is greater than 2.0,
they will form an ionic bond.
________ 15. Dispersion forces are weaker than hydrogen bonds.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
a. a substance in which all of the atoms are covalently
bonded to each other
________ 17. polar covalent bond
b. a bond formed when the atoms in a molecule are alike
and the bonding electrons are shared equally
________ 18. polar molecule
c. a term used to describe the weakest intermolecular
attractions; these include dispersion forces and dipole
interactions
________ 19. van der Waals forces
d. a bond formed when two different atoms are joined by a
covalent bond and the bonding electrons are shared
unequally
________ 20. network solid
e. a molecule in which one end is slightly positive and the
other end is slightly negative
188
Core Teaching Resources
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________ 16. nonpolar covalent
bond
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Name ___________________________
Date ___________________
Class __________________
Part D Questions and Problems
Answer the following in the space provided.
21. Arrange the following intermolecular attractions in order of increasing strength:
dipole interactions, dispersion forces, and hydrogen bonds.
22. State whether the following compounds contain polar covalent bonds, nonpolar covalent bonds, or ionic bonds, based on their electronegativities.
a.
b. SO2
b.
c. NO2
c.
d. Cl2
d.
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a. KF
Chapter 8 Covalent Bonding
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Date ___________________
Class __________________
COVALENT BONDING
Practice Problems
In your notebook, solve the following problems.
SECTION 8.1 MOLECULAR COMPOUNDS
1. Classify each of the following as an atom or a molecule.
a. Be
c. N2
b. CO2
d. H2O
e. Ne
2. Which of the following are diatomic molecules?
a. CO2
c. O2
b. N2
d. H2O
e. CO
3. What types of elements tend to combine to form molecular
compounds?
4. What information does a molecule’s molecular structure give?
5. How do ionic compounds and molecular compounds differ in their
relative melting and boiling points?
SECTION 8.2 THE NATURE OF COVALENT BONDING
2. Draw the electron dot structure for phosphorus trifluoride, PF3.
3. Draw the electron dot structure for nitrogen trichloride, NCl3.
4. Draw the electron dot configuration for acetylene, C2H2.
5. How many resonance structures can be drawn for CO32!? Show the
electron dot structures for each.
SECTION 8.3 BONDING THEORIES
1. Predict the shape and bond angle for the compound carbon
tetrafluoride, CF4.
2. Predict the shape and bond angle for phosphorus trifluoride, PF3.
3. Predict the type of hybridized orbitals involved in the compound boron
trichloride, BCl3.
4. What types of hybrid orbitals are involved in the bonding of the silicon
atoms in silicon tetrafluoride, SiF4?
5. Predict the shape and bond angle of fluorine monoxide, F2O.
190
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1. Draw the electron dot structure for hydrogen fluoride, HF.
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Date ___________________
Class __________________
6. Predict the shape of the CH2CF2 molecule. What hybridization is involved
in the carbon-carbon bonds?
7. How many sigma and pi bonds are used by each of the carbon atoms in
the following compound?
H O
H C1 C2 O H
H
SECTION 8.4 POLAR BONDS AND MOLECULES
1. What type of bond—nonpolar covalent, polar covalent, or ionic—will form
between each pair of atoms?
a. Na and O
b. O and O
c. P and O
2. Explain why most chemical bonds would be classified as either polar
covalent or ionic.
3. Would you expect carbon monoxide and carbon dioxide to be polar or
nonpolar molecules?
4. Draw the structural formulas for each molecule and identify polar covalent
bonds by assigning the slightly positive (!") and slightly negative (!#)
symbols to the appropriate atoms.
a. NH3
b. CF3
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5. Which would you expect to have the higher melting point, CaO or CS2?
Chapter 8 Covalent Bonding
191
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Name ___________________________
9.1
Date ___________________
Class __________________
NAMING IONS
Section Review
Objectives
• Determine the charges of monatomic ions by using the periodic table and write
the names of the ions
• Define a polyatomic ion and write the names and formulas of the most common
polyatomic ions
• Identify the two common endings for the names of most polyatomic ions.
Vocabulary
• monatomic ions
• polyatomic ions
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1
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Ions that consist of a single atom are called
Metallic elements tend to
3
2
1.
ions.
electrons. Group 1A ions have a
4
charge, whereas Group 2A metals form ions with a
charge, and Group 3A metals form ions with a
5
charge.
6
from the group number. For example, the
Group 7A elements form ions with a charge of
8
Many of the
10
9
system
11
The names of most common polyatomic ions end in either
12
or
13
.
6.
8.
9.
10.
naming system.
Ions containing more than one atom are called
4.
7.
.
have more than one common ionic
charge. These ions are named using either the
or the
7
3.
5.
The charge of a Group A nonmetal ion is determined by
subtracting
2.
ions.
11.
12.
13.
Chapter 9 Chemical Names and Formulas
211
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Name ___________________________
Date ___________________
Class __________________
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 14. The names of polyatomic ions end in -ite or -ate.
________ 15. In polyatomic ions for which there is an -ite/-ate pair, the -ite ending will
always indicate one less oxygen atom than the -ate ending.
________ 16. Polyatomic ions are anions.
________ 17. The charge on Group A metal ions is determined by subtracting the
group number from 8.
________ 18. The Group 6A ions have a charge of 2!.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
a. negatively charged ions
________ 20. polyatomic ions
b. ions formed from single atoms
________ 21. cations
c. a traditional way of naming transition metal cations
________ 22. anions
d. positively charged ions
________ 23. classical naming system
e. ions formed from groups of atoms
Part D Questions and Problems
Answer the following in the space provided.
24. What is the charge on a typical ion for each of the following groups?
a. 1A
c. 7A
b. 6A
d. 2A
25. Write the name of each of the following polyatomic ions.
a. HCO3!
c. MnO4!
b. NH4"
d. OH!
26. How many electrons does the neutral atom gain or lose to form each of the
following ions?
212
a. Ca2"
c. I!
b. S2!
d. Mn3"
Core Teaching Resources
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________ 19. monatomic ions
05_CTR_ch09 7/9/04 3:29 PM Page 213
Name ___________________________
Date ___________________
Class __________________
NAMING AND WRITING FORMULAS
FOR IONIC COMPOUNDS
9.2
Section Review
Objectives
• Apply the rules for naming and writing formulas for binary ionic compounds
• Apply the rules for naming and writing formulas for compounds with
polyatomic ions
Vocabulary
• binary compound
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1.
Binary ionic compounds are named by writing the name of
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the
1
followed by the name of the
binary compounds end in
3
2
2.
. Names of
4
. For example, NaI is
.
5
When a cation has more than one ionic charge, a
4.
5.
is used in the name.
Compounds with polyatomic ions whose names end in -ite
or -ate contain a polyatomic
6
that includes
7
.
In writing the formula of an ionic compound, the net ionic charge
must be
8
3.
6.
7.
8.
.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 9. The systematic name for baking soda (NaHCO3) is sodium
bicarbonate.
________ 10. In writing a formula for an ionic compound, the net ionic charge of
the formula must be zero.
Chapter 9 Chemical Names and Formulas
213
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Name ___________________________
Date ___________________
Class __________________
________ 11. Anions that contain oxygen end in -ite or -ate.
________ 12. The cation name is placed first when naming ionic compounds.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
________ 13. binary compounds
a. ions that consist of a single atom
________ 14. monatomic ions
b. ionic compounds composed of two elements
________ 15. polyatomic ions
c. Group B metals, many of which have more than one
common ionic charge
________ 16. transition metals
d. ions that consist of more than one atom
Part D Questions and Problems
Answer the following in the space provided.
17. Name the following compounds and tell what type of compound they are
(binary ionic or ionic with a polyatomic ion).
a. FeBr3
b. KOH
18. Write the formulas for the following compounds.
a. sodium chlorate
b. lead(II) phosphate
c. magnesium hydrogen carbonate
214
Core Teaching Resources
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c. Na2Cr2O7
05_CTR_ch09 7/9/04 3:29 PM Page 215
Name ___________________________
9.3
Date ___________________
Class __________________
NAMING AND WRITING FORMULAS
FOR MOLECULAR COMPOUNDS
Section Review
Objectives
• Interpret the prefixes in the names of molecular compounds in terms of their
chemical formulas
• Apply the rules for naming and writing formulas for binary molecular compounds
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1
Binary molecular compounds are composed of two
2
elements. The name of this type of compound ends in
Prefixes are used to show how many
3
of each element
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are present in a molecule of the compound. For example,
the name of As2S5 is
4
1.
.
2.
3.
4.
.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 5. Binary molecular compounds contain carbon.
________ 6. Charges must be balanced when writing formulas for molecular
compounds.
________ 7. CO2 is named monocarbon dioxide.
Chapter 9 Chemical Names and Formulas
215
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Name ___________________________
Date ___________________
Class __________________
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
________ 8. binary molecular
compound
a. used to indicate the relative number of atoms of an
element in a molecular compound
________ 9. prefix
b. prefix indicating one atom of an element in a molecule
________ 10. mono-
c. prefix indicating four atoms of an element in a molecule
________ 11. tetra-
d. nonionic compound containing atoms of two elements
Part D Questions and Problems
Answer the following in the space provided.
12. Name each of the following compounds.
a. PCl5
b. SO2
c. P4S10
13. Write formulas for the following compounds.
a. carbon tetrabromide
216
Core Teaching Resources
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b. dinitrogen tetroxide
05_CTR_ch09 7/9/04 3:29 PM Page 217
Name ___________________________
9.4
Date ___________________
Class __________________
NAMING AND WRITING FORMULAS
FOR ACIDS AND BASES
Section Review
Objectives
• Apply three rules for naming acids
• Apply the rules in reverse to write formulas of acids
• Apply the rules for naming bases
Vocabulary
• acid
• base
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number.
1
© Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
An acid is a compound that contains one or more
2
atoms and produces
1.
2.
when dissolved in water. There
3.
are rules for naming acids. For example, HBr is called
3
acid, whereas HNO3 is called
A base is a(n)
5
4
4.
acid.
6
compound that produces
when
dissolved in water. Ionic compounds that are bases are named
in the same way as other
of the
8
7
compounds, that is, the name
is followed by the name of the
9
.
5.
6.
7.
8.
9.
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 10. A compound that contains hydrogen atoms will be an acid when
dissolved in water.
Chapter 9 Chemical Names and Formulas
217
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Name ___________________________
Date ___________________
Class __________________
________ 11. An acid contains one or more hydroxide ions.
________ 12. Chemists have a special system for naming bases.
Part C Matching
Match each description in Column B to the correct term in Column A.
Column A
Column B
________ 13. acid
a. a compound containing hydrogen that ionizes to yield
hydrogen ions in solution
________ 14. base
b. a solution in which the solvent is water
________ 15. aqueous solution
c. a compound that produces hydroxide ions in water
Part D Questions and Problems
Answer the following in the space provided.
16. Write the formula for each acid or base.
a. magnesium hydroxide
b. hydrofluoric acid
d. lithium hydroxide
17. Name each acid or base.
a. KOH
b. HI
c. H2SO4
218
Core Teaching Resources
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c. phosphoric acid
05_CTR_ch09 7/9/04 3:29 PM Page 219
Name ___________________________
9.5
Date ___________________
Class __________________
THE LAWS GOVERNING FORMULAS
AND NAMES
Section Review
Objectives
• Define the laws of definite proportions and multiple proportions
• Apply the rules for writing chemical formulas by using a flowchart
• Apply the rules for naming chemical compounds by using a flowchart
Vocabulary
• law of definite proportions
• law of multiple proportions
Part A Completion
Use this completion exercise to check your understanding of the concepts and terms
that are introduced in this section. Each blank can be completed with a term, short
phrase, or number. [Use Figure 9.20 to complete this exercise.]
© Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
1
states that in samples of any chemical
1.
compound, the masses of the elements are always in the same
2.
The law of
2
3
. The law of
states that whenever the same two
3.
elements form more than one compound, the different masses
4.
of one element that combine with the same mass of the other
5.
4
6.
element are in the ratio of
H3PO4 is a(n)
5
numbers.
6
. It is called
7
7.
.
. It contains two elements, so it is a
8.
8
compound. It does not contain a metal, so it is a binary
9.
9
compound. The compound is called
CCl4 is not a(n)
Pb(C2H3O2)2 is not a(n)
12
11
10.
.
. It contains more than two
13
. C2H3O2! is a polyatomic
metal. The compound is called
10
15
.
. Pb is a Group
14
11.
12.
13.
14.
15.
Chapter 9 Chemical Names and Formulas
219
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Name ___________________________
Date ___________________
Class __________________
Part B True-False
Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT.
________ 16. Roman numerals are used when naming Group B metal cations.
________ 17. Names of compounds containing polyatomic anions end in -ide.
________ 18. Prefixes are used when naming binary ionic compounds.
________ 19. Compounds containing two elements are called binary compounds.
Part C Questions and Problems
Answer the following in the space provided.
20. Name the following compounds.
a. Pb(C2H3O2)4
b. HF
c. P2O5
d. LiBr
21. Write formulas for the following compounds.
a. phosphorus pentachloride
b. iron(II) oxide
d. potassium chloride
e. calcium nitrate
220
Core Teaching Resources
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c. nitric acid
05_CTR_ch09 7/9/04 3:29 PM Page 221
Name ___________________________
9
Date ___________________
Class __________________
CHEMICAL NAMES AND FORMULAS
Practice Problems
In your notebook, solve the following problems.
SECTION 9.1 NAMING IONS
1. What is the charge on the ion typically formed by each element?
a. oxygen
c. sodium
e. nickel, 2 electrons lost
b. iodine
d. aluminum
f. magnesium
2. How many electrons does the neutral atom gain or lose when each ion
forms?
a. Cr3!
c. Li!
e. Cl"
b. P3"
d. Ca2!
f. O2"
3. Name each ion. Identify each as a cation or an anion.
a. Sn2!
c. Br"
e. H"
b. Co3!
d. K!
f. Mn2!
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4. Write the formula (including charge) for each ion. Use Table 9.3 if necessary.
a. carbonate ion
c. sulfate ion
e. chromate ion
b. nitrite ion
d. hydroxide ion
f. ammonium ion
5. Name the following ions. Identify each as a cation or an anion.
a. CN"
c. PO43"
e. Ca2!
b. HCO3"
d. Cl"
f. SO32"
SECTION 9.2 NAMING AND WRITING FORMULAS
FOR IONIC COMPOUNDS
1. Write the formulas for these binary ionic compounds.
a. magnesium oxide
c. potassium iodide
e. sodium sulfide
b. tin(II) fluoride
d. aluminum chloride
f. ferric bromide
2. Write the formulas for the compounds formed from these pairs of ions.
a. Ba2!, Cl"
c. Ca2!, S2"
e. Al3!, O2"
b. Ag!, I"
d. K!, Br"
f. Fe2!, O2"
3. Name the following binary ionic compounds.
a. MnO2
c. CaCl2
e. NiCl2
g. CuCl2
b. Li3N
d. SrBr2
f. K2S
h. SnCl4
Chapter 9 Chemical Names and Formulas
221
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Name ___________________________
Date ___________________
Class __________________
4. Write formulas for the following ionic compounds.
a. sodium phosphate
c. sodium hydroxide
e. ammonium chloride
b. magnesium sulfate
d. potassium cyanide
f. potassium dichromate
5. Write formulas for compounds formed from these pairs of ions.
a. NH4!, SO42"
c. barium ion and hydroxide ion
b. K!, NO3"
d. lithium ion and carbonate ion
6. Name the following compounds.
a. NaCN
c. Na2SO4
e. Cu(OH)2
b. FeCl3
d. K2CO3
f. LiNO3
7. Name and give the charge of the metal cation in each of the following
ionic compounds.
a. Na3PO4
c. CaS
e. FeCl3
b. NiCl2
d. K2S
f. CuI
SECTION 9.3 NAMING AND WRITING FORMULAS
FOR MOLECULAR COMPOUNDS
1. Name the following molecular compounds.
a. PCl5
c. NO2
e. P4O6
g. SiO2
b. CCl4
d. N2F2
f. XeF2
h. Cl2O7
2. Write the formulas for the following binary molecular compounds.
c. sulfur dioxide
b. dichlorine monoxide
d. dinitrogen tetrafluoride
SECTION 9.4 NAMING AND WRITING FORMULAS
FOR ACIDS AND BASES
1. Name the following compounds as acids.
a. HNO2
b. H2SO4
c. HF
d. H2CO3
2. Write the formulas for the following bases.
a. calcium hydroxide
c. aluminum hydroxide
b. ammonium hydroxide
d. lithium hydroxide
SECTION 9.5 THE LAWS GOVERNING FORMULAS AND NAMES
1. Write the formulas for these compounds.
222
a. potassium sulfide
e. hydrobromic acid
i. sulfur hexafluoride
b. tin(IV) chloride
f. aluminum fluoride
j. magnesium chloride
c. hydrosulfuric acid
g. dinitrogen pentoxide
k. phosphoric acid
d. calcium oxide
h. iron(III) carbonate
l. nitric acid
Core Teaching Resources
© Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
a. nitrogen tribromide
05_CTR_ch09 7/9/04 3:29 PM Page 223
Name ___________________________
Date ___________________
Class __________________
2. Complete this table by writing correct formulas for the compounds formed by
combining positive and negative ions.
SO42!
NO3!
PO43!
OH!
Ca2"
Al3"
Na"
Pb4"
3. Name the following compounds.
a. K3PO4
c. NaHSO4
e. N2O5
g. PI3
b. Al(OH)3
d. HgO
f. NBr3
h. (NH4)2SO4
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4. Explain the difference between the law of definite proportions and the law of
multiple proportions.
Chapter 9 Chemical Names and Formulas
223