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CHAPTER 2
Atoms, Molecules, and Ions
CHAPTER TERMS AND DEFINITIONS
Numbers in parentheses after definitions give the text sections in which the terms are explained. Starred
terms are italicized in the text. Where a term does not fall directly under a text section heading,
additional information is given for you to locate it.
atomic theory explanation of the structure of matter in terms of different combinations of very
small particles (2.1)
atom minute particle of which matter is composed; the smallest part of an element that can enter
into chemical reaction (2.1)
element
substance whose atoms all have the same atomic number (2.1, 2.3)
compound type of matter composed of atoms of two or more elements chemically combined in
fixed proportions (2.1)
chemical reaction rearrangement of atoms present in reacting substances to give new chemical
combinations present in the substances formed (2.1)
atomic symbol
element (2.1)
one- or two-letter notation used to represent an atom corresponding to a particular
law of multiple proportions
when two elements form more than one compound, the masses of one
element in these compounds for a fixed mass of the other element are in ratios of small whole numbers
(2.1)
nucleus
positively charged central core of an atom; contains most of the atom’s mass (2.2)
electron very light, negatively charged particle that exists in the region around the positively
charged nucleus (2.2)
cathode*
anode*
negative electrode (2.2)
positive electrode (2.2)
cathode rays*
(2.2)
rays that originate from the cathode, or negative electrode, in a gas-discharge tube
coulomb (C)*
unit of electric charge (2.2)
nuclear model* most of the mass of an atom is concentrated in a positively charged center, called
the nucleus, around which negatively charged electrons move (2.2)
proton nuclear particle having a positive charge equal to +e (e being the charge on an electron) and
a mass more than 1800 times that of an electron (2.3)
atomic number (Z)
neutron
number of protons in an atomic nucleus; identifies the element (2.3)
neutral particle of mass almost identical to that of a proton but without electric charge (2.3)
mass number (A)
total number of protons and neutrons in a nucleus (2.3)
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Chapter 2: Atoms, Molecules, and Ions
nuclide
an atom characterized by a definite atomic number and mass number (2.3)
nuclide symbol* symbol for a nuclide in which the mass number is written as a superscript and the
atomic number as a subscript on the left of the symbol for the element (2.3)
isotopes
atoms whose nuclei have the same atomic number but different mass numbers (2.3)
mass spectrometer*
atomic mass
units (2.4)
instrument used to determine atomic mass (2.4)
average atomic mass for the naturally occurring element, expressed in atomic mass
atomic mass unit (amu)
mass unit equal to exactly one-twelfth the mass of a carbon-12 atom (2.4)
mass spectrum* chart recording from the mass spectrometer that shows the relative numbers of
atoms for various masses (2.4)
fractional abundance
(2.4)
fraction of the total number of atoms that is composed of a particular isotope
periodic table tabular arrangement of elements in rows and columns, highlighting the regular
repetition of properties of the elements (2.5)
period (of periodic table)
elements in any one horizontal row of the periodic table (2.5)
group (of periodic table)
elements in any one column of the periodic table (2.5)
main-group (representative) elements*
transition elements*
elements in the B groups of the periodic table (2.5)
inner-transition elements*
lanthanides*
actinides*
two rows of elements at the bottom of the periodic table (2.5)
first of the two rows of inner-transition elements (2.5)
second of the two rows of inner-transition elements (2.5)
alkali metals*
halogens*
elements in the A groups of the periodic table (2.5)
elements in Group IA of the periodic table (2.5)
elements in Group VIIA of the periodic table (2.5)
metal substance or mixture that has a characteristic luster or shine and is generally a good conductor
of heat and electricity; elemental metals are to the left of the staircase line on the periodic table (2.5)
malleable*
ductile*
able to be hammered into sheets (2.5)
able to be drawn into wire (2.5)
nonmetal element to the right of the staircase line on the periodic table; exhibits characteristics
different from those of metals (2.5)
metalloid (semimetal) element bordering the staircase line on the periodic table; exhibits both
metallic and nonmetallic properties (2.5)
semiconductors* elements that, when pure, are poor conductors of electricity at room temperature
but become good conductors at higher temperatures (2.5)
doping* adding small amounts of other elements to pure semiconductor elements to make them very
good electrical conductors (2.5, margin note)
chemical formula notation that uses atomic symbols with numerical subscripts to convey the
relative proportions of atoms of the different elements in the substance (2.6)
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Chapter 2: Atoms, Molecules, and Ions
23
molecule
definite group of atoms that are chemically bonded together and, as a group, electrically
neutral (2.6)
molecular substance*
substance composed of molecules all of which are alike (2.6)
molecular formula
gives the exact number of different atoms of an element in a molecule (2.6)
structural formula*
molecule (2.6)
chemical formula that shows which atoms are bonded to one another in a
polymers very large molecules that are made up of a number of smaller molecules repeatedly linked
together (2.6)
monomers
the small molecules that are linked together to form a polymer (2.6)
ion electrically charged particle obtained from an atom or chemically bonded group of atoms by
addition or removal of one or more electrons (2.6)
anion
negatively charged ion (2.6)
cation
positively charged ion (2.6)
ionic compound
crystal*
(2.6)
compound composed of cations and anions (2.6)
solid having a regular three-dimensional arrangement of either ions, atoms, or molecules
formula unit
group of atoms or ions explicitly symbolized in the chemical formula (2.6)
organic compounds molecular substances that contain carbon combined with other elements, such
as hydrogen, oxygen, and nitrogen (2.7)
hydrocarbons
compounds containing only hydrogen and carbon (2.7)
functional groups
alcohol*
ether*
reactive portion of a molecule that undergoes predictable reactions (2.7)
molecule that contains an –OH functional group (2.7)
organic molecule that contains an oxygen atom between two carbon atoms (2.7)
chemical nomenclature
inorganic compounds
monatomic ion
systematic naming of chemical compounds (2.8)
compounds composed of elements other than carbon (2.8)
ion formed from a single atom (2.8)
Stock system* system for naming compounds in which a Roman numeral within parentheses
follows the first-named element to indicate its charge or oxidation number (2.8)
oxidation state* or oxidation number*
hypothetical charge assigned in accordance with certain
rules; denoted with a Roman numeral following the name of the metal atom (2.8, margin note)
-ous* in an older system of nomenclature, a suffix added to the stem name of an element to indicate
the cation of lower charge; also indicates the oxoacid with fewer oxygen atoms (2.8)
-ic* in an older system of nomenclature, a suffix added to the stem name of an element to indicate
the cation of higher charge; indicates the oxoacid with more oxygen atoms; also indicates an acid
solution obtained from binary compounds of hydrogen and nonmetals (2.8)
-ide* suffix added to the stem name of the element to name monatomic anions or the more
electronegative element in binary compounds (2.8)
polyatomic ion ion consisting of two or more atoms chemically bonded together and carrying a net
electric charge (2.8)
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Chapter 2: Atoms, Molecules, and Ions
oxoanion (oxyanion)*
element (2.8)
anion composed of oxygen with another element, which is the central
acid* molecular compound that can yield one or more hydronium ions, H3O+, and an anion for each
acid molecule when the acid dissolves in water (2.8)
oxoacid
acid containing hydrogen, oxygen, and another element (2.8)
-ate*
suffix denoting the oxoanion with the greater number of oxygen atoms (2.8)
-ite*
suffix denoting the oxoanion with the lesser number of oxygen atoms (2.8)
hypo-*
(2.8)
prefix denoting the oxoacid or oxoanion with the least number of oxygen atoms in the series
per-* prefix denoting the oxoacid or oxoanion with the greatest number of oxygen atoms in the
series (2.8)
acid anions*
di-*
anions that have hydrogen atoms they can lose as hydronium ions, H3O+ (2.8)
Greek prefix meaning two (2.8)
thio-*
prefix meaning an oxygen in the root ion name has been replaced by a sulfur atom (2.8)
binary compound
compound composed of only two elements (2.8)
hydro-* prefix added to the stem name of a nonmetal to name the acid solution obtained from
binary compounds of hydrogen and nonmetals (2.8)
hydrate
compound that contains water molecules weakly bound in its crystals (2.8)
chemical equation
(2.9)
symbolic representation of a chemical reaction in terms of chemical formulas
reactant starting substance in a chemical reaction; appears to the left of the arrow in a chemical
equation (2.9)
product substance that results from a chemical reaction; appears to the right of the arrow in a
chemical equation (2.9)
coefficient* number that appears in front of a formula in a chemical equation and gives the relative
number of molecules or formula units of a substance involved in the reaction (2.9)
(g)*
(2.9)
phase label placed after a formula in a chemical equation to indicate that the substance is a gas
(l)* phase label placed after a formula in a chemical equation to indicate that the substance is a
liquid (2.9)
(s)*
(2.9)
phase label placed after a formula in a chemical equation to indicate that the substance is a solid
(aq)* phase label placed after a formula in a chemical equation to indicate that the substance is in
aqueous (water) solution (2.9)
catalyst*
balanced*
substance that speeds up a reaction without undergoing any net change itself (2.9)
describes a chemical equation having correct coefficients (2.10)
balancing by inspection* trial-and-error method of balancing a chemical equation by writing
appropriate coefficients until there is the same number of any one elemental atom on each side of the
arrow (2.10)
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Chapter 2: Atoms, Molecules, and Ions
25
CHAPTER DIAGNOSTIC TEST
1.
Dalton’s atomic theory postulated that matter
a.
is in continuous motion.
b.
is continuous (infinitely divisible) in nature.
c.
changes in mass when heated to combustion.
d.
can exist in three states—gas, liquid, and solid.
e.
is composed of small particles called atoms.
2.
Robert Millikan’s _______________ experiment, in conjunction with Thomson’s value for m/e of
the electron, allowed an accurate calculation of the mass of the electron.
3.
In a mass spectrograph, the natural isotopes of iron were observed to have the following atomic
masses (and percentage abundances): 53.94 (5.84%), 55.94 (91.68%), 57.94 (2.17%), and 57.93
(0.310%). From these data, give the average atomic mass of iron.
4.
Explain the difference in meaning between the symbols H, H2, and H2O.
5.
Complete the following statements about aluminum nitrate, Al(NO3)3.
a.
One formula unit of aluminum nitrate contains _____ Al atom(s).
b.
One formula unit of aluminum nitrate contains _____ N atom(s).
c.
One formula unit of aluminum nitrate contains _____ O atom(s).
6.
A nucleus consists of 32 protons and 41 neutrons. What is the nuclide symbol for this nucleus?
7.
In a sample of hydrogen gas there are 4.92  1018 hydrogen atoms. How many methane molecules
(formula CH4) could be formed?
8.
When the ions Na+ and CO32 combine chemically, the compound sodium carbonate (called soda
ash) is formed. Keeping in mind that chemical formulas are written as electrically neutral species,
write the correct formula for sodium carbonate.
9.
Write the molecular formulas for the molecules having the following structural formulas.
a.
H
H
H
C
C
H
C
H
H
b.
H
H
C
H
H
N
Cl
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26
Chapter 2: Atoms, Molecules, and Ions
c.
F
F
N
N
F
d.
F
F—O—O—F
10. Match each term in the left-hand column with a descriptive example in the right-hand column.
___
1.
compound
a.
Sn
___
2.
Rutherford
b.
characterized by Z and A
___
3.
atomic symbol
c.
1/12 the mass of a C-12 atom
___
4.
J. J. Thomson
d.
gold-foil experiment
___
5.
nuclide
e.
cathode-ray-tube experiments
___
6.
amu
f.
methane (CH4)
11. On the following diagram of the periodic table, indicate the regions where nonmetals, metals, and
metalloids are to be found.
12. Fluorine, chlorine, and bromine are _______________ (metals/metalloids/nonmetals) that belong
to the Group VIIA family, commonly referred to as the ____________________.
13. Carbonic acid, H2CO3, exists only in aqueous solution and is formed when CO2 dissolves in water.
This is what gives carbonated drinks their “sparkling” taste. Give the formula and name of the
oxoanion of carbonic acid.
14. Tell whether you would expect the following compounds to be ionic or molecular in nature.
a.
XeF4
b.
CS2
c.
NaI
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Chapter 2: Atoms, Molecules, and Ions
27
15. Complete the following chart with the appropriate numbers or symbols.
Symbol
C
Al3+
?
?
?
Atomic number
?
?
1
?
47
Protons
?
?
?
16
?
Neutrons
?
14
?
?
61
Mass number
14
?
1
32
?
Electrons
?
?
1
18
46
16. Write the formula for each of the following substances.
a.
zinc hydrogen carbonate
b.
copper(II) dichromate dihydrate
c.
manganous hydroxide
d.
bismuth nitride
e.
plumbous iodide
f.
periodic acid
17. Write the name of each of the following substances. Where appropriate, give the Stock system
name and the common name.
a.
Cr(NO3)3 ∙ 9H2O
b.
AlPO4
c.
Fe2S3
d.
PbSO3
e.
Co(CN)2
f.
P2O5
g.
CO
18. Fill in the blanks, referring to the following:
 3Fe + 4H2O
Fe3O4 + 4H2 
The preceding expression is called a chemical (a) ______. It describes a chemical (b) ______ in
which a chemical (c) ______ occurs in the identity of the reacting molecules. Write a description
of the information it gives you: (d) ______.
ANSWERS TO CHAPTER DIAGNOSTIC TEST
If you missed an answer, study the text section and problem-solving skill given (PS Sk.) in parentheses
after the answer.
1.
e (2.1)
2.
oil-drop (2.2)
3.
55.9 amu (2.4, PS Sk. 2)
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28
4.
Chapter 2: Atoms, Molecules, and Ions
H is the symbol for the element hydrogen. It represents one atom of hydrogen. H2 represents a
molecule of the element hydrogen, which is made up of two atoms of hydrogen. H2O represents a
molecule of a compound (two or more different elements combined together) that contains two
atoms of hydrogen and one atom of oxygen. (2.6)
5.
a.
1
b.
3
c.
9 (2.6)
6.
73
32 Ge
7.
1.23  1018 molecules (2.6)
8.
Na2CO3 (2.6, PS Sk. 3)
(2.3, PS Sk. 1)
9.
a.
C3H6
b.
CH4NCl
c.
N2F4
d.
O2F2 (2.6)
10. (1) f (2.1), (2) d (2.2), (3) a (2.1), (4) e (2.2), (5) b (2.3), (6) c (2.4)
11.
Nonmetals
M
Metals
et
al
lo
id
s
(2.5)
12. nonmetals, halogens (2.5)
13. CO32– is the carbonate ion. (2.8, PS Sk. 6)
14.
a.
molecular
b.
molecular
c.
ionic (2.8, PS Sk. 5)
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Chapter 2: Atoms, Molecules, and Ions
29
15.
Symbol
Atomic
number
Protons
Neutrons
Mass number
Electrons
C
6
Al3+
13
H
1
S2
16
Ag+
47
6
8
14
6
13
14
27
10
1
0
1
1
16
16
32
18
47
61
108
46
(2.3, PS Sk. 1)
16.
a.
Zn(HCO3)2
b.
CuCr2O7 ∙ 2H2O
c.
Mn(OH)2
d.
BiN
e.
PbI2
f.
HIO4 (2.8, PS Sk. 4)
a.
chromium(III) nitrate nonahydrate
b.
aluminum phosphate
c.
iron(III) sulfide or ferric sulfide
d.
lead(II) sulfite or plumbous sulfite
e.
cobalt(II) cyanide or cobaltous cyanide
f.
diphosphorus pentoxide
g.
carbon monoxide (2.8, PS Sk. 4)
a.
equation
b.
reaction
c.
change
d.
1 formula unit of Fe3O4 (iron oxide) reacts with 4 molecules of H2 (hydrogen) to form 3
atoms of Fe (iron) and 4 molecules of H2O (water). (2.9, 2.10)
17.
18.
SUMMARY OF CHAPTER TOPICS
The definitions presented in this chapter are central to the language of chemistry. Work at mastering
them as soon as possible. The descriptions of the terms atom, element, and compound are theoretical
explanations that form the basis of our understanding of chemistry. Besides knowing their definitions, it
is equally important that you know how atoms, elements, and compounds behave in laboratory work.
Both theoretical and practical descriptions are given in the list of chapter terms and definitions.
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30
Chapter 2: Atoms, Molecules, and Ions
To keep ion names and charges on the tip of your tongue, use text Tables 2.4 and 2.5 and study guide
Table 2.1 to make flip cards on 3 × 5 in index cards. Write the name of the ion or element on one side
and the symbol (with charge, if an ion) on the other side. Flip through the cards, putting the ones you
don’t know in a separate pile. Work on the ones you don’t know in your spare time.
2.1 Atomic Theory of Matter
Learning Objectives

List the postulates of atomic theory.

Define element, compound, and chemical reaction in the context of these postulates.

Recognize the atomic symbols of the elements.

Explain the significance of the law of multiple proportions.
2.2 The Structure of the Atom
Learning Objectives

Describe Thomson’s experiment in which he discovered the electron.

Describe Rutherford’s experiment that led to the nuclear model of the atom.
In the cathode-ray tube, the electrode through which the electrons enter the tube, the cathode, is
designated the negative electrode. The electrode at which the electrons leave the tube, the anode, is
electrically positive with respect to the cathode. In Chapter 18 you will learn more about electricity as it
relates to chemistry.
J. J. Thomson established the mass-to-charge ratio of the electron, and Robert Millikan determined the
electric charge on an electron. Shortly thereafter, Ernest Rutherford, on the basis of his gold-foil
experiment, postulated a nuclear model of the atom.
2.3 Nuclear Structure; Isotopes
Learning Objectives

Name and describe the nuclear particles making up the nucleus of the atom.

Define atomic number, mass number, and nuclide.

Write the nuclide symbol for a given nuclide.

Define and provide examples of isotopes of an element.

Write the nuclide symbol of an element. (Example 2.1)
Problem-Solving Skill
1.
Writing nuclide symbols. Given the number of protons and neutrons in a nucleus, write its
nuclide symbol (Example 2.1).
The atomic number (the number of protons in the nucleus) tells what element we are dealing with.
Every calcium atom has 20 protons in its nucleus. The calcium atom loses two electrons to become the
Ca2+ ion. This charged atom still has 20 protons in the nucleus. The bromine atom gains one electron to
become the bromide ion, Br, with 36 electrons. This charged atom still has 35 protons in its nucleus.
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Chapter 2: Atoms, Molecules, and Ions
31
Exercise 2.1
A nucleus consists of 17 protons and 18 neutrons. What is its nuclide symbol?
Known : The complete symbol includes the symbol for the element, the mass number, and the
atomic number. Atomic number = number of protons; mass number = number of protons +
number of neutrons. (See table on inside back cover of the text.)
Solution: Atomic number = 17; mass number = 17 + 18 = 35. The symbol is
35
17 Cl.
The atomic numbers and average atomic masses of the elements are given in the periodic table on
the inside front cover of the text. Each element is represented by a square with the atomic symbol
in it. The number above the atomic symbol is the atomic number.
2.4 Atomic Masses
Learning Objectives

Define atomic mass unit and atomic mass.

Describe how a mass spectrometer can be used to determine the fractional abundance of the
isotopes of an element.

Determine the atomic mass of an element from the isotopic masses and fractional
abundances. (Example 2.2)
Problem-Solving Skill
2.
Determining atomic mass from isotopic masses and fractional abundances.
Given the isotopic masses (in atomic mass units) and fractional isotopic abundances for a naturally
occurring element, calculate its atomic mass (Example 2.2).
If you look at text Table 2.1, you will see that the relative masses of protons and neutrons are not
exactly 1 amu. Thus the mass number of an isotope (the sum of the number of protons and neutrons)
always will be a whole number, but the mass of the isotope (in amu) will not be. You will notice from
Example 2.2 that we can get the mass number by rounding off the isotopic mass to a whole number.
To take an average, we usually add the given values and divide by the number of them. For instance,
the average of 5, 7, and 9 is 5 + 7 + 9 = 2 1/3 = 7. This method gives us the correct answer because each
value has equal weight or representation. We could have gotten the same average, 7, by taking ⅓ of
each value and then adding them:
1
1
1
7
9
5
21
(5) + (7) + (9) =
+ 3 + 3 = 3 =7
3
3
3
3
This second method exemplifies a weighted average. We must use this method for determining average
relative atomic masses because there is never an equal number of atoms of each isotope in any naturally
occurring sample of an element. We could give the amount of each isotope as a percentage of the total,
but it is more useful to give the value in decimal form, which we call the fractional abundance. For
example, in Example 2.2 the fractional abundance of chromium-50 is 0.0435. This means that 4.35% of
the atoms in any naturally occurring sample of chromium are the chromium-50 isotope. You should
memorize this method of finding atomic masses.
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Chapter 2: Atoms, Molecules, and Ions
Exercise 2.2
Chlorine consists of the following isotopes:
Isotope
Chlorine-35
Chlorine-37
Mass (amu)
34.96885
36.96590
Fractional Abundance
0.75771
0.24229
What is the atomic mass of chlorine?
Known:
The atomic mass is the weighted average of isotopic masses.
Solution: 34.96885 amu  0.75771 = 26.4962 amu
36.96590 amu  0.24229 = 8.95647 amu
Atomic mass of chlorine = 35.453 amu
Although the actual masses of atoms are known, the relative atomic masses (called atomic
masses) are much easier to use in calculations. For example, the actual average atomic mass of a
calcium atom is 6.656  1023 g, whereas its relative average atomic mass (atomic mass) is 40.08
amu. Be sure you understand the difference between the two and the relationship between them.
2.5 Periodic Table of the Elements
Learning Objectives

Identify periods and groups on the periodic table.

Find the main-group and transition elements on the periodic table.

Locate the alkali-metal and halogen groups on the periodic table.

Recognize the portions of the periodic table that contain the metals, nonmetals, and
metalloids (semimetals).
There are many periodic phenomena in our experience. The phases of the moon are periodic; every 28
days we see a full moon. The seasons of the year are periodic; there is a regular repetition of winter,
spring, summer, and fall.
Exercise 2.3
By referring to the periodic table (Figure 2.15 or inside front cover of the text), identify the group and
period to which each of the following elements belongs. Then decide whether the element is a metal,
nonmetal, or metalloid.
a.
Se
b.
Cs
c.
Fe
d.
Cu
e.
Br
Known: Elements to the left of the periodic table staircase line with characteristic properties are
metals; elements to the right of the line with characteristic properties are nonmetals; elements
bordering the line with properties of metals and nonmetals are metalloids.
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Chapter 2: Atoms, Molecules, and Ions
33
Solution:
a.
Selenium is a nonmetal. It is to the right of the staircase line, in Period 4 and Group VIA.
b.
Cesium is a metal. It is to the left of the line, in Period 6 and Group IA, the alkali metals.
c.
Iron is a metal. It is in Period 4 and Group VIIIB.
d.
Copper is a metal. It is in Period 4 and Group IB.
e.
Bromine is a nonmetal. It is in Period 4 and Group VIIA, the halogens.
2.6 Chemical Formulas; Molecular and Ionic Substances
Learning Objectives

Determine when the chemical formula of a compound represents a molecule.

Determine whether a chemical formula is also a molecular formula.

Define ion, cation, and anion.

Classify compounds as ionic or molecular.

Define and provide examples for the term formula unit.

Specify the charge on all substances, ionic and molecular.

Write an ionic formula, given the ions. (Example 2.3)
Problem-Solving Skill
3.
Writing an ionic formula, given the ions. Given the formulas of a cation and an anion, write the
formula of the ionic compound of these ions (Example 2.3).
Exercise 2.4
Potassium chromate is an important compound of chromium (see Figure 2.23). It is composed of K+
and CrO42 ions. Write the formula of the compound.
Known:
Compounds must be electrically neutral.
Solution: There must be two K+ ions to provide two positive charges to neutralize the 2 charge
on the chromate ion. The formula is K2CrO4.
2.7 Organic Compounds
Learning Objectives

List the attributes of molecular substances that make them organic compounds.

Explain what makes a molecule a hydrocarbon.

Recognize some functional groups of organic molecules.
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34
Chapter 2: Atoms, Molecules, and Ions
2.8 Naming Simple Compounds
Learning Objectives

Recognize inorganic compounds.

Learn the rules for predicting the charges of monatomic ions in ionic compounds.

Apply the rules for naming monatomic ions.

Learn the names and charges of common polyatomic ions.

Name an ionic compound from its formula. (Example 2.4)

Write the formula of an ionic compound from its name. (Example 2.5)

Determine the order of elements in a binary (molecular) compound.

Learn the rules for naming binary molecular compounds, including the Greek prefixes.

Name a binary compound from its formula. (Example 2.6)

Write the formula of a binary compound from its name. (Example 2.7)

Name a binary molecular compound from its molecular model. (Example 2.8)

Recognize molecular compounds that are acids.

Determine whether an acid is an oxoacid.

Learn the approach for naming binary acids and oxoacids.

Write the name and formula of an anion from the acid. (Example 2.9)

Recognize compounds that are hydrates.

Learn the rules for naming hydrates.

Name a hydrate from its formula. (Example 2.10)

Write a formula of a hydrate from its name. (Example 2.11)
Problem-Solving Skills
4.
Writing the name of a compound from its formula, or vice versa. Given the formula of a
simple compound (ionic, binary molecular, acid, or hydrate), write the name (Examples 2.4, 2.6,
and 2.10), or vice versa (Examples 2.5, 2.7, and 2.11).
5.
Writing the name and formula of an anion from the acid. Given the name and formula of an
oxoacid, write the name and formula of the oxoanion, or given the name and formula of an
oxoanion, write the formula and name of the oxoacid (Example 2.9).
A Chemist Looks at: Thirty Seconds on the Island of Stability
Questions for Study
1.
Which elements are unstable and fall apart by radioactive decay?
2.
What is the heaviest naturally occurring element?
3.
How are the transuranium elements generally made in the laboratory?
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Chapter 2: Atoms, Molecules, and Ions
35
4.
How are elements represented in nuclear equations?
5.
What would you predict to be the nature of the yet-to-be-discovered elements 113, 115, and 117?
Answers to Questions for Study
1.
Those elements with nuclides having atomic numbers greater than 83 are unstable and undergo
radioactive decay.
2.
The heaviest naturally occurring element is uranium (element 92), which is unstable and
undergoes radioactive decay.
3.
The transuranium elements are made by bombarding elements of fairly high atom mass with
particles of lower mass.
4.
In nuclear equations, the elements are represented using their nuclide symbols, which include
information on the atomic mass number (number of protons and neutrons) and atomic number
(number of protons).
5.
Elements 113, 115, and 117 would be seventh-row elements in the periodic table and should
exhibit properties consistent with Groups 3A, 5A, and 7A, respectively.
You must learn the element names and symbols and the ion names, formulas, and charges to master
naming and writing the formulas of compounds. Refer to text Tables 2.4 and 2.5 or to study guide Table
2.1. Make flip cards to help you.
It also will be useful for you to know the formulas and names of common acids. They are given in text
Table 2.7 and study guide Table 2.2.
Exercise 2.5
Write the names of the following compounds.
a.
CaO
b.
PbCrO4
a.
Known: CaO is a binary compound and will end in -ide. Ca (Group IIA) forms the 2+ ion;
thus the oxygen ion is O2–.
Solution: Calcium oxide
b.
Known: The chromate ion has a charge of 2–. The lead ion is thus Pb2+.
Solution: Lead(II) chromate, or plumbous chromate
Exercise 2.6
A compound has the name thallium(III) nitrate. What is the formula? (The symbol of thallium is Tl.)
Known:
Tl (Group IIIA) has a charge of 3+; nitrate is NO3; the compound must be neutral.
Solution: Tl(NO3)3
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36
Chapter 2: Atoms, Molecules, and Ions
Table 2.1 Charges, Formulas, and Names of Some Common Ions
Cations
Charge Formula Name
1+
NH4+
ammonium
Cs+
cesium
+
Cu
copper(I); cuprous
H3O+
hydronium
+
Li
lithium
K+
potassium
Rb+
rubidium
+
Ag
silver
Na+
sodium
2+
Ba2+
Be2+
Cd2+
Ca2+
Co2+
Cu2+
Fe2+
Pb2+
Mg2+
Mn2+
Hg22+
2+
Hg2+
Ni2+
Sr2+
Sn2+
Zn2+
3+
4+
Al3+
Bi3+
Cr3+
barium
beryllium
cadmium
calcium
cobalt(II); cobaltous
copper(II); cupric
iron(II); ferrous
lead(II); plumbous
magnesium
manganese(II);
manganous
mercury(I);
mercurous
mercury(II);
mercuric
nickel(II); nickelous
strontium
tin(II);stannous
zinc
Fe3+
aluminum
bismuth
chromium(III);
chromic
iron(III); ferric
Sn4+
tin(IV); stannic
Anions
Charge Formula
1
C2H3O2
Br
Cl
ClO
ClO2
ClO3
ClO4
CN
H2PO4
F
H
HCO3
OH
I
NO3
NO2
MnO4
Name
acetate
bromide
chloride
hypochlorite
chlorite
chlorate
perchlorate
cyanide
dihydrogen phosphate
fluoride
hydride
hydrogen carbonate;
bicarbonate
hydrogen sulfate;
bisulfate
hydrogen sulfite;
bisulfite
hydroxide
Iodide
nitrate
nitrite
permanganate
CO32
carbonate
C2O42
CrO42
Cr2O72
HPO42
O2
O22
SO42
S2
SO32
S2O32
oxalate
chromate
dichromate
monohydrogen phosphate
oxide
peroxide
sulfate
sulfide
sulfite
thiosulfate
AsO43
N3
PO43
arsenate
nitride
phosphate
HSO4
HSO3
2
3
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Chapter 2: Atoms, Molecules, and Ions
37
Table 2.2 Formulas and Names of Some Common Acids
Formula
HF(aq)
HCl(aq)
HBr(aq)
HI(aq)
H2S(aq)
HCN(aq)
HC2H3O2
H2CO3
HClO
Name
hydrofluoric acid
hydrochloric acid
hydrobromic acid
hydroiodic acid
hydrosulfuric acid
hydrocyanic acid
acetic acid
carbonic acid
hypochlorous acid
Formula
HClO2
HClO3
HClO4
HNO3
HNO2
H2SO4
H2SO3
H3AsO4
H3PO4
Name
chlorous acid
chloric acid
perchloric acid
nitric acid
nitrous acid
sulfuric acid
sulfurous acid
arsenic acid
phosphoric acid
Following are practice grids and answers for writing chemical formulas and names for ionic and
covalent compounds.
Formula/Nomenclature Practice Grid for Ionic Compounds
Anion
Cation
Sodium
Fluoride
Chloride Oxide
Sulfide
Sulfate
Sulfite
Carbonate Chlorate
Potassium
Magnesium
Calcium
Aluminum
Chromium(II)
Iron(II)
Iron(III)
Copper(I)
Copper(II)
Silver
Ammonium
This grid is designed to help you learn to name a compound when the formula is given and to write the
formula when the name of the compound is given. To use this grid effectively:
1.
read aloud the name of the compound as you write the formula, and
2.
once all formulas are written, cover up the names of the cations and anions and recite the names of
the compounds from the formulas.
Note: This is a naming exercise; some of the named compounds do not exist.
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38
Chapter 2: Atoms, Molecules, and Ions
Formula/Nomenclature Practice Grid for Molecular Compounds
Fluorine
Formula Name
Chlorine
Formula Name
Oxygen
Formula Name
Sulfur
Formula Name
Boron(III)
Silicon(IV)
Phosphorus(V)
Tungsten(VI)
This grid is designed to help you learn to name molecular compounds. To use this grid effectively:
1.
write the formula for the compound formed between the element in the far left column (which is
written as if it were an ion) and the nonmetal element listed across the top of the grid, and
2.
write the name using Greek number prefixes.
Copyright © Houghton Mifflin Company. All rights reserved.
Chapter 2: Atoms, Molecules, and Ions
Answers
Formula/Nomenclature Practice Grid for Ionic Compounds
Anion
Cation
Fluoride
Chloride
Oxide
Sulfide
Sulfate
Sulfite
Carbonate
Chlorate
Sodium
NaF
NaCl
Na2O
Na2S
Na2SO4
Na2SO3
Na2CO3
NaClO3
Potassium
KF
KCl
K2O
K2S
K2SO4
K2SO3
K2CO3
KClO3
Magnesium
MgF2
MgCl2
MgO
MgS
MgSO4
MgSO3
MgCO3
Mg(ClO3)2
Calcium
CaF2
CaCl2
CaO
CaS
CaSO4
CaSO3
CaCO3
Ca(ClO3)2
Aluminum
AlF3
AlCl3
Al2O3
Al2S3
Al2(SO4)3
Al2(SO3)3
Al2(CO3)3
Al(ClO3)3
Chromium(II)
CrF2
CrCl2
CrO
CrS
CrSO4
CrSO3
CrCO3
Cr(ClO3)2
Iron(II)
FeF2
FeCl2
FeO
FeS
FeSO4
FeSO3
FeCO3
Fe(ClO3)2
Iron(III)
FeF3
FeCl3
Fe2O3
Fe2S3
Fe2(SO4)3
Fe2(SO3)3
Fe2(CO3)3
Fe(ClO3)3
Copper(I)
CuF
CuCl
Cu2O
Cu2S
Cu2SO4
Cu2SO3
Cu2CO3
CuClO3
Copper(II)
CuF2
CuCl2
CuO
CuS
CuSO4
CuSO3
CuCO3
Cu(ClO3)2
Silver
AgF
AgCl
Ag2O
Ag2S
Ag2SO4
Ag2SO3
Ag2CO3
AgClO3
Ammonium
NH4F
NH4Cl
(NH4)2O
(NH4)2S
(NH4)2SO4
(NH4)2SO3
(NH4)2CO3
NH4ClO3
Copyright © Houghton Mifflin Company. All rights reserved.
39
40
Chapter 2: Atoms, Molecules, and Ions
Answers
Formula/Nomenclature Practice Grid for Molecular Compounds
Fluorine
Formula Name
Chlorine
Formula Name
Formula
Oxygen
Name
Formula
Sulfur
Name
Boron(III)
BF3
boron
trifluoride
BCl3
boron
trichloride
B2O3
diboron
trioxide
B2S3
diboron
trisulfide
Silicon(IV)
SiF4
silicon
tetrafluoride
SiCl4
silicon
tetrachloride
SiO2
silicon
dioxide
SiS2
silicon
disulfide
Phosphorus(V)
PF5
phosphorus
pentafluoride
PCl5
phosphorus
pentachloride
P2O5
diphosphorus
pentoxide
P2S5
diphosphorus
pentasulfide
Tungsten(VI)
WF6
tungsten
hexafluoride
WCl6
tungsten
hexachloride
WO3
tungsten
trioxide
WS3
tungsten
trisulfide
Exercise 2.7
Name the following compounds: (a) Cl2O6, (b) PCl3, (c) PCl5.
Known: For binary compounds, we name the first element and add a suffix to the stem name of
the second element. For elements forming more than one compound, we use the Greek prefixes
listed in text Table 2.6, omitting mono- for a first-named element.
Solution: (a) The text list of compounds in which the Greek prefixes are used shows that chlorine
and oxygen form more than one compound. The name would be dichlorine hexoxide. (b)
phosphorus trichloride (c)phosphorus pentachloride
Exercise 2.8
Give formulas for the following compounds.
a.
carbon disulfide
b.
sulfur trioxide
Known: We change the names of the elements to symbols and translate the prefixes to subscripts.
Solution:
a.
CS2
b.
SO3
Exercise 2.9
Using the molecular models, name the following chemical compounds: (a) BF3 (b) H2Se.
Known: Because these are binary molecular compounds, we use Greek prefixes listed in text
Table 2.6 and name the elements in the order in which they appear in the chemical formula.
Solution: (a) boron trifluoride (b) dihydrogen selenide
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Chapter 2: Atoms, Molecules, and Ions
41
Exercise 2.10
What is the name and formula of the anion corresponding to perbromic acid, HBrO4?
Known: We remove H+ from the acid to get the ion. We name the ion by replacing -ic in the
acid name with -ate.
Solution: The ion is BrO4–, and it is the perbromate ion.
Exercise 2.11
Washing soda has the formula Na2CO3 ∙ 10H2O. What is the chemical name of this substance?
Known: The formula indicates that the compound is a hydrate. We name hydrates from the
anhydrous compound, followed by the word hydrate with a Greek prefix to indicate the associated
number of water molecules.
Solution: The name is sodium carbonate decahydrate.
Exercise 2.12
Photographers’ hypo, used to fix negatives during the development process, is sodium thiosulfate
pentahydrate. What is the chemical formula of this compound?
Known: The anhydrous compound is composed of sodium ions (Na+) and thiosulfate ions
(S2O32); pentahydrate means there are five associated molecules of water.
Solution: The formula is Na2S2O3 ∙ 5H2O.
2.9 Writing Chemical Equations
Learning Objectives

Identify the reactants and products in a chemical equation.

Write chemical equations using appropriate phase labels, symbols of reaction conditions,
and the presence of a catalyst.
2.10 Balancing Chemical Equations
Learning Objectives

Determine if a chemical reaction is balanced.

Master techniques for balancing chemical equations. (Example 2.12)
Problem-Solving Skill
6.
Balancing simple equations. Given the formulas of the reactants and products in a chemical
reaction, obtain the coefficients of the balanced equation (Example 2.12).
Balancing equations does not have to be a difficult process. The principle behind it is the law of
conservation of matter. Usually, balancing equations is a trial-and-error process. Always check your
final answer to be sure that you have the same number of atoms of each element on each side of the
arrow.
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42
Chapter 2: Atoms, Molecules, and Ions
Exercise 2.13
Find the coefficients that balance the following equations.
a.
 POCl3
O2 + PCl3 
b.
 P4O6 + N2
P4 + N2O 
c.
 As2O3 + SO2
As2S3 + O2 
d.
 Ca(H2PO4)2
Ca3(PO4)2 + H3PO4 
Solution:
a.
 2POCl3
O2 + 2PCl3 
b.
 P4O6 + 6N2
P4 + 6N2O 
c.
 2As2O3 + 6SO2
2As2S3 + 9O2 
d.
 3Ca(H2PO4)2
Ca3(PO4)2 + 4H3PO4 
ADDITIONAL PROBLEMS
1.
Supply the missing information in the following table.
Symbol
Atomic number
Number of protons
Number of neutrons
Number of electrons
Mass number
C
?
?
8
?
?
Al3+
?
?
?
?
27
?
?
16
16
18
?
2.
Explain the nature of isotopes of elements in terms of the subatomic particles of matter.
3.
Write the symbol for
a.
a nucleus containing 8 protons and 8 neutrons.
b.
the carbon-14 nucleus.
4.
A sample of neon always contains the three isotopes of neon: neon-20, neon-21, and neon-22. The
natural abundances of these isotopes are 90.92%, 0.257%, and 8.82%, respectively. Their isotopic
masses are 19.99244, 20.99395, and 21.99138 amu, respectively. Calculate the atomic mass of
neon.
5.
Write the formula for the compound of each of the following ion pairs.
a.
Ca2+ and H–
b.
Al3+ and OH–
c.
Mg2+ and PO43–
d.
NH4+ and SO42–
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Chapter 2: Atoms, Molecules, and Ions
43
6.
Classify each of the following elements as a metal, a metalloid, or a nonmetal: S, Sr, Co, N, Ga, I,
Ar, Si, Li, V.
7.
Complete the following chart with the appropriate numbers or symbols.
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
8.
9.
P
?
?
?
31
?
?
11
?
12
?
10
H
?
?
2
?
?
Mg2+
?
?
?
24
?
?
29
?
?
64
29
Write the formula for each of the following substances.
a.
sodium sulfite heptahydrate
b.
iron(III) chloride
c.
calcium fluoride
d.
sulfuric acid
e.
silicon tetrabromide
f.
barium hydrogen sulfite
g.
perchloric acid
h.
dinitrogen tetroxide
Name each of the following compounds. Where appropriate, give both the Stock system name and
the common name.
a.
SrBr2 ∙ 6H2O
b.
MnCl2
c.
HBrO3
d.
NO2
e.
LiH
f.
BCl3
10. For each of the following, write the symbol and name for the corresponding oxoacid or oxoanion.
a.
HClO, hypochlorous acid
b.
PO43–, phosphate ion
c.
HNO3, nitric acid
d.
SO32–, sulfite ion
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44
Chapter 2: Atoms, Molecules, and Ions
11. Determine which of the following statements is (are) incorrect with respect to the following
chemical equation:
 2HCl
H2 + Cl2 
a.
Two atoms of hydrogen occur on both sides of the equation.
b.
The chemical reaction described by this equation consists of a rearrangement of the atoms of
hydrogen and chlorine to give hydrogen chloride (HCl).
c.
Since one molecule of H2 and one molecule of Cl2 react, two molecules of product must
form.
d.
This equation describes a chemical change.
12. Balance each of the following equations.
a.
 ___ Na2O
___ Na + ___ O2 
b.
 ___ AsH3
___ As + ___ H2 
c.
 ___ Ba(NO3)2 + ___ H2O
___ Ba(OH)2 + ___ HNO3 
d.
 ___ Al2(SO4)3 + ___ H2
___ Al + ___ H2SO4 
e.
 ___ CO2 + H2O
___ C3H8 + ___ O2 
13. A solution of lead sulfate reacts with a solution of sodium chloride to form solid lead chloride and
aqueous sodium sulfate. Write the balanced equation for this reaction, including state (phase)
labels.
ANSWERS TO ADDITIONAL PROBLEMS
If you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses
after the answer.
1.
Symbol
Atomic number
Number of protons
Number of neutrons
Number of electrons
Mass number
C
6
6
8
6
14
S2–
16
16
16
18
32
Al3+
13
14
10
27
(2.3)
2.
Isotopes are atoms of the same element that have the same number of protons and the same
number of electrons but different numbers of neutrons. Therefore, isotopes have different masses.
(2.3)
3.
a.
16
8O
b.
14
6C
(2.3, PS Sk. 1)
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Chapter 2: Atoms, Molecules, and Ions
4.
The atomic mass is the weighted average of isotopic masses.
19.99244 amu  0.9092
=
18.1771 amu
20.99395 amu  0.00257
=
0.0540 amu
21.99138 amu  0.0882
=
1.9396 amu
Atomic mass of neon
=
20.1707 = 20.17 amu
(2.4, PS Sk. 2)
5.
6.
a.
CaH2
b.
Al(OH)3
c.
Mg3(PO4)2
d.
(NH4)2SO4 (2.6, PS Sk. 3)
The metals are Sr, Co, Ga, Li, and V. Si is a metalloid. The nonmetals are S, N, I, and Ar. (2.5)
7.
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
P
15
15
16
31
15
Na+
11
11
12
23
10
H
1
1
2
3
1
Mg2+
12
12
12
24
10
Cu
29
29
35
64
29
(2.3)
8.
a.
Na2SO3 ∙ 7H2O
b.
FeCl3
c.
CaF2
d.
H2SO4
e.
SiBr4
f.
Ba(HSO3)2
g.
HClO4
h.
N2O4 (2.8, PS Sk. 4)
a.
strontium bromide hexahydrate
b.
manganese(II) chloride, manganous chloride
c.
bromic acid
d.
nitrogen dioxide
e.
lithium hydride
f.
boron trichloride (2.8, PS Sk. 4)
9.
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45
46
Chapter 2: Atoms, Molecules, and Ions
10.
a.
ClO–, hypochlorite ion
b.
H3PO4, phosphoric acid
c.
NO3–, nitrate ion
d.
H2SO3, sulfurous acid (2.8, PS Sk. 6)
11. c (2.10)
12.
a.
 2Na2O
4Na + O2 
b.
 2AsH3
2As + 3H2 
c.
 Ba(NO3)2 + 2H2O
Ba(OH)2 + 2HNO3 
d.
 Al2(SO4)3 + 3H2
2Al + 3H2SO4 
e.
 3CO2 + 4H2O (2.10, PS Sk. 7)
C3H8 + 5O2 
 PbCl2(s) + Na2SO4(aq) (2.6, 2.8, 2.9, 2.10, PS Sk. 6, 7)
13. PbSO4(aq) + 2NaCl(aq) 
CHAPTER POST-TEST
1.
Indicate whether each of the following statements is true or false. If a statement is false, change it
so that it is true.
a.
According to Dalton’s atomic theory, compounds are kinds of matter composed of atoms of
two or more elements. True/False: ______________________________________________
__________________________________________________________________________
b.
A chemical symbol is used to designate elements, and formulas are used to designate
formula units. True/False: _____________________________________________________
__________________________________________________________________________
c.
A molecular formula does not indicate the arrangement of atoms in a molecule.
True/False:_________________________________________________________________
__________________________________________________________________________
d.
KClO3, Na2S, and BF3 are ionic compounds. True/False: ____________________________
__________________________________________________________________________
2.
If you have one dozen formula units of Mg(OH)2, how many hydrogen atoms do you have?
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Chapter 2: Atoms, Molecules, and Ions
3.
47
The structure of sulfuric acid, the most widely used chemical, is
O
H
O
S
O
H
O
Write the molecular formula for sulfuric acid. (Hint: The order of atoms is H, S, and then O.)
4.
5.
6.
Which one of the following statements correctly describes the difference between a structural and
a molecular formula?
a.
The structural formula indicates the composition of the molecule and the spatial arrangement
of the atoms, and the molecular formula shows how the different atoms bond with each
other.
b.
The structural formula represents the simplest composition of a molecule, and the molecular
formula represents the actual composition.
c.
The structural formula shows how the atoms are bonded together in a molecule, and the
molecular formula shows the atomic composition of the molecule.
d.
None of the above are correct statements.
Write the symbol and name for the corresponding oxoacid or oxoanion.
a.
BrO2–, bromite ion
b.
HNO2, nitrous acid
c.
IO3–, iodate ion
d.
C2H3O2–, acetate ion
Complete the following chart with the appropriate numbers or symbols.
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
Cr
?
?
28
?
?
?
26
?
?
56
24
Br
?
?
?
80
?
?
?
92
143
?
92
?
24
?
29
?
24
7.
Classify each of the following elements as a metal, a nonmetal, or a metalloid: Mg, Cu, Pb, As,
Cl.
8.
Which of the following is an incorrect statement? Rewrite that statement so that it is correct.
a.
Na and Cs are in the same group in the periodic table.
b.
The nonmetallic elements are on the far right side of the periodic table.
c.
Most of the known chemical elements are classified as metals.
d.
Elements in the same period of the periodic table have similar properties.
e.
Fluorine is classified as a nonmetallic element.
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48
9.
Chapter 2: Atoms, Molecules, and Ions
Write the formula for each of the following.
a.
barium peroxide
b.
silver chlorite
c.
strontium oxalate monohydrate
d.
ammonium iodate
e.
stannous nitrite
f.
phosphoric acid
10. Write the name of each of the following. Where appropriate, give the Stock system name and the
common name.
a.
KClO4
b.
Hg2Br2
c.
Mg(C2H3O2)2 ∙ 4H2O
d.
CaH2
e.
NaMnO4
f.
HNO3
11. Fill in the blanks, referring to the following:

2ZnS(s) + 3O2(g)  2SO2(g) + 2ZnO(s)
This equation is a symbolic representation of the (a) _____ between two (b) _____ of (c) _____ in
the (d) _____ state and three (e) _____ of oxygen (f) _____ to produce two (g) _____ of gaseous
(h) _____ and two formula units of solid (i) _____. The triangle over the arrow indicates that the
reactants are (j) _____ to produce the products.
12. Balance the following equations.
a.
 ___ P4O10
___ P + ___ O2 
b.
 ___ Ag2SO4 + ___ Ni
___ Ag + ___ NiSO4 
c.
 ___ CO2 + ___ H2O
___ C6H14 + ___ O2 
d.
 ___ SiF4 + ___ H2O
___ HF + ___ SiO2 
e.
 ___ Na2CO3 + ___ H2O + ___ CO2
___ NaHCO3 
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Chapter 2: Atoms, Molecules, and Ions
49
ANSWERS TO CHAPTER POST-TEST
If you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses
after the answer.
1.
a.
True. (2.1)
b.
True. (2.1, 2.6)
c.
True. (2.6)
d.
False. KClO3 and Na2S are ionic compounds; BF3 is a molecular compound. (2.8)
2.
two dozen (2.1)
3.
H2SO4 (2.6)
4.
c (2.6)
5.
a.
HBrO2, bromous acid
b.
NO2–, nitrite ion
c.
HIO3, iodic acid
d.
HC2H3O2, acetic acid (2.8, PS Sk. 5)
6.
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
Cr
24
24
28
52
24
Fe2+
26
26
30
56
24
Br
35
35
45
80
36
U
92
92
143
235
92
Cr
24
24
29
53
24
(2.3)
7.
Mg, Cu, and Pb are metals, Cl is a nonmetal, and As is a metalloid. (2.5)
8.
d. Elements in the same group of the periodic table have similar properties. (2.5)
9.
a.
BaO2
b.
AgClO2
c.
SrC2O4 ∙ H2O
d.
NH4IO3
e.
Sn(NO2)2
f.
H3PO4 (2.8, PS Sk. 4)
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50
Chapter 2: Atoms, Molecules, and Ions
10.
a.
potassium perchlorate
b.
mercury(I) bromide or mercurous bromide
c.
magnesium acetate tetrahydrate
d.
calcium hydride
e.
sodium permanganate
f.
nitric acid (2.8, PS Sk. 4)
a.
chemical reaction
b.
formula units
c.
zinc sulfide
d.
solid (2.9, 2.10)
e.
molecules
f.
gas
g.
molecules
h.
sulfur dioxide
i.
zinc oxide
j.
heated
a.
 P4O10
4P + 5O2 
b.
 Ag2SO4 + Ni
2Ag + NiSO4 
c.
 12CO2 + 14H2O
2C6H14 + 19O2 
d.
 SiF4 + 2H2O
4HF + SiO2 
e.
 Na2CO3 + H2O + CO2 (2.10, PS Sk. 6)
2NaHCO3 
11.
12.
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