Chapter 3. Elements, Atoms, Ions, and the Periodic Table
3.1 The Periodic Law and the Periodic Table
Mendeleev, Dmitri (1834-1907), Russian chemist is best
known for his work on crating a periodic table for 63 then
known elements.
The periodic law is an organized "map" of the
elements that relates their structure to their chemical and
physical properties. The periodic table is the result of the
periodic law, and provides the basis for prediction of such
properties as relative atomic and ionic size, ionization energy,
and electron affinity, as well as metallic or non-metallic
character and reactivity.
The modern periodic table exists in several forms.
The most important variation is in group numbering. The tables in the text use the two
most commonly accepted numbering systems.
Numbering Groups in the Periodic Table
Periods and Groups
Periods are the horizontal rows of elements in the periodic table; the columns
represent groups or families.
The lanthanide series and the actinide series are parts of periods 6 and 7, respectively,
and groups that have been named include the alkali metals, the alkaline earth metals, the
halogens, and the noble gases. Group A elements are called representative elements;
Group B elements are transition elements. Metals, metalloids, and nonmetals can be
identified by their location on the periodic table.
These groups are number from 1 - 18, left to right and groups have their Roman
numbers and A or B classification..
Group IA Group IIA -
alkali metal: Li, Na, K Rb, Cs, Fr
alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra
Group IIIA - B, Al, Ga, In, Tl
Group IVA - C, Si, Ge, Sn, Pb
Group VA - N, P, As, Sb, Bi
Group VIA - O, S, Se, Te, Po
Group VIIA - Halogens: Cl, Br, I, At
Group VIIIA - Noble gases: He, Ne, Ar, Kr, Xe, Rn
In addition to groups in the periodic table there are three blocks of elements called
transition elements which are labeled with B, Lanthanides and Actinides
Metals, Nonmetals and Metalloids
Most of the elements in the periodic table are metals. They are found to the left of the
table. In between metal and non-metals there are semi-metals or metalloids.
Atomic Number and Atomic Mass
Atomic number (Z) is shown on top of the space for each element. At the bottom
average atomic mass calculated based on isotopes of each elements is written.
The atomic number of an element represents the number of protons in the nucleus of
atoms of that specific element. No other element has that specific number of protons.
Atomic numbers can be found on most Periodic Charts/Tables. The atomic number will
always be a whole number value without decimals. Look on a periodic chart at the
elements listed below. Do you know how to find an elements atomic number?
Use your periodic table to find the atomic number and atomic mass rounded to two
decimal place of each of the following elements:
a) Hydrogen (H) b) Krypton (Kr) c) Potassium (K) d) Gold (Au)
3.2 Electron Arrangement and the Periodic Table
Valence Electrons
Bohr concluded that the energy levels of an atom can handle only a certain number of
electrons at a time.
The Quantum Mechanical Atom
Energy Levels and Sublevels
Building Atoms by Orbital Filling
Schrodinger's work showed that each orbital could have a maximum of two electrons.
Energy levels could contain different numbers of orbitals.
Energy levels further from the nucleus can accommodate more orbitals than
energy levels nearer the nucleus.
Energy levels can have sublevels when multiple orbitals are present.
Consider the energy levels and their orbitals by consulting Fig 3.5 [DTC064].
maximum number of electrons per level
Energy Levels
number of sublevels
Sublevels
2
n=1
one
1s
8
n=2
two
2s 2p
18
n=3
three
3s 3p 3d
32
n=4
four
4s 4p 4d 4f
maximum number of electrons per sublevel
2
2 + 6 2 + 6 + 10
2 + 6 + 10 + 14
Images of Orbital
s orbital
p orbital
d orbital
f orbital
image of three p orbitals occupying the x, y, and z
axes around the nucleu
Consider this diagram which shows the
relationships of the energy sublevels.
This is the same concept as Fig 3.6 [DTC066] in
your book.
The lowest energy sublevels always fill first.
Note how Level 3 and Level 4 sublevels overlap
in energy.
This would be grusome to memorize.
Can you believe that the The
Periodic Table displays all the data
in this figure?
Amazingly, the "electron configurations" of the elements are "embedded" in the Periodic Table.
Honk, if you can see this "embedded" information in the Periodic Table?
Analogy: The periodic table is actually a packing slip that tells how the electrons are packed around the
nucleus.
The maximum number of electrons that can be in an energy level is 2n 2, where
n is equal to the energy level being considered.
Energy
Level
n=1
n=2
n=3
n=4
maximum number of
electrons
in an Energy Level
2
8
18
32
# of
sublevels
Sublevels names
maximum number of electrons
per sublevel
1
2
3
4
2
2, 6
2, 6, 10
2, 6, 10, 14
s
s, p
s, p, d
s, p, d, f
Electron Configuration and the Aufbau Principle
The outermost electrons in an atom are valence electrons. For representative elements, the number of
valence electrons in an atom corresponds to the group or family number. Metals tend to have fewer
valence electrons than nonmetals.
Electron configuration of the elements is predictable, using the aufbau principle. Knowing the
electron configuration, we can identify valence electrons and begin to predict the kinds of reactions that
the elements will undergo.
Elements in the last family, the noble gases, have either two or eight valence electrons. Their most
important properties are their extreme stability and lack of reactivity. A full energy level is responsible
for this unique stability.
Abbreviated Electron Configurations
3.3 The Octet Rule
The octet rule tells us that in chemical reactions elements will gain, lose or share the
minimum number of electrons necessary to achieve the electron configuration of the
nearest inert gas.
Metallic elements tend to form cations and nonmetals form anions that are
isoelectronic with their nearest noble gas neighbor.
Ion Formation and the Octet Rule
3.4 Trends in the Periodic Table
Atomic size increases from top to bottom but decreases from left to right in the periodic
table. Cations are smaller than the parent atom. Anions are larger than the parent atom.
Ions with multiple positive charge are even smaller than their corresponding
monopositive ion; ions with multiple negative charge are larger than their
corresponding less negative ion.
The energy required to remove an electron from the atom is the ionization energy.
Descending a group, the ionization energy decreases. Proceeding across a period, the
ionization energy increases.
The energy released when a single electron is added to neutral atom in the gaseous
state is known as the electron affinity. Electron affinities generally decrease proceeding
down a group and increase proceeding across a period.
Exceptions exist for periodic trends. They are generally small anomalies, and do not
detract from the predictive power of the periodic table.
Atomic Size
Ion Size
Ionization Energy
Electron Affinity
Instructional Objectives
Conceptual Objectives
 Know the meaning and ramifications of the periodic law.
 Recognize the important subdivisions of the periodic table, periods, groups (families), metals, and
nonmetals.
 Understand the relationship between the electronic structure of an element and its position in the
periodic table.
 Know the meaning of the octet rule and its predictive usefulness.
 Understand the meaning of and utility of ionization energies and electron affinities in predicting ion
formation.
Performance Objectives
 Extract information about an element from the periodic table, for example, the mass, number of
protons, neutrons, and electrons in an atom of any element.
 Write electron configurations for atoms of the most commonly occurring elements.
 Use the octet rule to predict the charge of common cations and anions.
 Use the periodic table and its predictive power to estimate relative size of atoms and ions, as well as
relative magnitudes of ionization energy and electron affinity.
Health Applications
 Recognize the importance of trace metals in the diet and the effect of deficiencies, such as Wilson’s
disease.
In-Chapter Examples
Example 3 1:
Example 3.2:
Determining electron arrangement.
Writing the electron configuration of tin.
Chapter Overview
The Periodic Law and the Periodic Table
The periodic law is an organized "map" of the elements that relates their structure to their chemical
and physical properties. The periodic table is the result of the periodic law, and provides the basis for
prediction of such properties as relative atomic and ionic size, ionization energy, and electron affinity, as
well as metallic or non-metallic character and reactivity.
The modern periodic table exists in several forms. The most important variation is in group
numbering. The tables in the text use the two most commonly accepted numbering systems.
Periods are the horizontal rows of elements in the periodic table; the columns represent groups or
families. The lanthanide series and the actinide series are parts of periods 6 and 7, respectively, and
groups that have been named include the alkali metals, the alkaline earth metals, the halogens, and the
noble gases. Group A elements are called representative elements; Group B elements are transition
elements. Metals, metalloids, and nonmetals can be identified by their location on the periodic table.
Electron Arrangement and the Periodic Table
The Octet Rule
Trends in The Periodic Table
Atomic size increases from top to bottom but decreases from left to right in the periodic table.
Cations are smaller than the parent atom. Anions are larger than the parent atom. Ions with multiple
positive charge are even smaller than their corresponding monopositive ion; ions with multiple negative
charge are larger than their corresponding less negative ion.
The energy required to remove an electron from the atom is the ionization energy. Descending a
group, the ionization energy decreases. Proceeding across a period, the ionization energy increases.
The energy released when a single electron is added to neutral atom in the gaseous state is known as
the electron affinity. Electron affinities generally decrease proceeding down a group and increase
proceeding across a period.
Exceptions exist for periodic trends. They are generally small anomalies, and do not detract from the
predictive power of the periodic table.
Hints for Faster Coverage
Although the discussion detailing the importance of ionization energy and electron affinity in ion
formation helps to give the student a more complete picture, it is not essential for the understanding of the
material which is presented in subsequent chapters. Consequently, it may be abbreviated or omitted.
Suggested Problem Sets
The Periodic Law and the Periodic Table - 15, 17, 19, 21, 23, 25, 27, 29
Electron Arrangement and the Periodic Table - 31, 33, 35, 37, 39, 41, 43, 45, 47, 49
The Octet Rule - 51, 53, 55, 57, 59
Trends in the Periodic Table - 61, 63, 65, 67, 69
In-Chapter Perspectives
A MEDICAL PERSPECTIVE: Copper Deficiency and Wilson's Disease. This perspective uses the
element copper to describe the biochemical importance of trace minerals in the diet. The role of dietary
copper in the function of the central nervous system is discussed, as well as the serious consequences of
copper deficiency.
A CLINICAL PERSPECTIVE: Dietary Calcium. This perspective describes the essential role that the
element calcium plays in our growth and development, as well as our long-term health. It was once
believed that our need for calcium diminished significantly once we passed adolescence. This is clearly
not true, based on research into the variety of bodily functions in which calcium plays a critical part.
Additional Perspectives
Other perspectives that can be inserted into the lectures dealing with periodic properties might
revolve around the special utility of certain elements. For example, silicon and germanium can be
discussed in relation to the semiconductor industry...the calculators which students use function because
the integrated circuits contain these elements. Also, the unique properties of the transition metals give rise
to some useful practical chemistry...metallic mercury, a liquid in the thermometers (thermometers
function because of the marked change in volume of mercury with change in temperature). Additional,
potentially useful topics include: copper is a good electrical conductor; many trace metals are essential
nutrients.
Critical Thinking Problems
1.
2.
3.
4.
5.
This problem will engage the student in the topic of the periodic table by relating practical, everyday
materials to the concept of elements and element classification.
The predictive nature of the periodic table is emphasized in this question. Also, the fact that the
periodic table is a "work in progress" illustrates the evolutionary character of science in general.
The relationship between chemical and physical properties and technological application is
emphasized. The student is encouraged to go beyond the textbook and do some library research.
The students will have an opportunity to use a thought process similar to that which Mendeleev
himself may have used. A number of logical tables can be constructed; a class discussion of the
reasons underlying the student's choice could be useful.
The octet rule works well for only certain elements, this exercise helps to explain why this is so.
Chapter 3
Elements, Atoms, Ions, and the Periodic Table
Solutions to the Even-Numbered Problems
In-Chapter Questions and Problems
3.2
a.
b.
c.
d.
Ar (argon)
C (carbon)
B (boron)
Ar (argon)
3.4
a. magnesium, atomic number = 12, mass = 24.31 amu
b. neon, atomic number = 10, mass = 20.18 amu
c. selenium, atomic number = 34, mass 78.96 amu
3.6
a.
b.
c.
d.
e.
Total electrons = 19, valence electrons = 1
Total electrons = 9, valence electrons = 7
Total electrons = 15, valence electrons = 5
Total electrons = 8, valence electrons = 6
Total electrons = 20, valence electrons = 2
3.8
a. Potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
b. Phosphorus: 1s2, 2s2, 2p6, 3s2, 3p3
3.10
a. [Ar] 4s1
b. [Ne] 3s2, 3p3
3.12
a.
b.
c.
d.
e.
3.14
a. (smallest) F, Cl, Br, I (largest)
b. (highest) F, Cl, Br, I (lowest)
c. (highest) Cl, F, Br, I (lowest)
Note: the exception to the general trend for chlorine and fluorine.
Cl– has 18 electrons, Ar has 18 electrons; isoelectronic
Na+ has 10 electrons, Ne has 10 electrons; isoelectronic
Mg2+ has 10 electrons, Na+ has 10 electrons; isoelectronic
Li+ has 2 electrons, Ne has 10 electrons; not isoelectronic
O2– has 10 electrons, F– has 10 electrons; isoelectronic
End-of-Chapter Questions and Problems
3.16
a. electronic configuration - the arrangement of electrons, in orbits or orbitals,
around a nucleus of an atom.
b. octet rule - the rule which predicts that atoms form the most stable molecules
or ions when they are surrounded by eight electrons in their highest occupied
energy level.
c. ionization energy - the energy required to remove an electron from an atom
in the gas phase.
d. isoelectronic - a property exhibited by two or more ions, atoms, or molecules
that have the same number of electrons.
3.18
a. false
b. true
3.20
a.
b.
c.
d.
3.22
a. calcium
The metals are: Ca, K, Cu, Zn
The representative metals are: Ca, K
The elements that tend to form positive ions are: Ca, K, Cu, Zn
The element that is inert is Kr
b. copper
c. cobalt
3.24
Group IIA is known collectively as the alkaline earth metals and consists of
beryllium, magnesium, calcium, strontium, barium, and radium.
3.26
Group VIIIA is known collectively as the noble gases and consists of helium,
neon, argon, krypton, xenon, and radon.
3.28
a. boron
b. silicon
c. arsenic
3.30
a. Metals are located to the left of the "staircase" division of the periodic table
and include such elements as sodium, calcium, and copper.
b. Metalloids are located along the "staircase" and include boron, silicon, and
arsenic.
c. Nonmetals are located to the right of the "staircase" and include such
elements as carbon, nitrogen, and fluorine.
3.32
a.
b.
c.
d.
e.
f.
3.34
The outermost energy level of the elements in group VIIIA is complete. Helium
has two outermost electrons; all of the other members of the family have a
complete octet of electrons.
3.36
a.
b.
c.
d.
3.38
A sublevel is a part of a principal energy level and is designated s, p, d, and f.
Each sublevel may contain one or more orbitals, regions of space containing a
maximum of two electrons with their spins paired.
3.40
A 2s orbital is a higher energy orbital than a 1s orbital because it is a part of a
higher energy principal energy level.
3.42
A sketch of the three p orbitals can be found in Figure 3.3, on page 63 of the
textbook.
px, py, and pz orbitals differ only in their special orientation, along the x, y, and z
axes of a set of x, y, z coordinates.
two
one
four
seven
zero (or eight)
zero (or eight)
n = 1; 2n2 = 2(1)2 = 2
n = 2; 2n2 = 2(2)2 = 8
n = 3; 2n2 = 2(3)2 = 18
n = 4; 2n2 = 2(4)2 = 32
3.44
An orbital can hold a maximum of two electrons.
3.46
a.
b.
c.
d.
two s electrons; 2 e–/orbital x 1 orbital = 2 e–
six p electrons; 2 e–/orbital x 3 orbitals = 6 e–
ten d electrons; 2 e–/orbital x 5 orbitals = 10 e–
fourteen f electrons; 2 e–/orbital x 7 orbital = 14 e–
3.48
a.
b.
c.
d.
e.
f.
B (5 e–) 1s2, 2s2, 2p1
S (16 e–) 1s2, 2s2, 2p6, 3s2, 3p4
Ar (18 e–) 1s2, 2s2, 2p6, 3s2, 3p6
V (23 e–) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d3
Cd (48 e–) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10
Te (52 e–) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p4
3.50
a., c., and d. are incorrect.
For a., 4 electrons, the element is helium, and the correct electron configuration
is:
1s2, 2s2
For c., 11 electrons, the element is sodium, and the correct electron
configuration is:
1s2, 2s2, 2p6, 3s1
For d., 5 electrons, the element is boron, and the correct electron configuration
is;
1s2, 2s2, 2p1
I– (54 e–)
Ba2+ (54 e–)
Se2– (36 e–)
Al3+ (10 e–)
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6
1s2, 2s2, 2p6
3.52
a.
b.
c.
d.
3.54
a. F–, 10e–; Cl–, 18 e–; Not isoelectronic
b. K+, 18 e–; Ar, 18e–; Isoelectronic
3.56
Noble gases are nonreactive because they all have a complete outer shell.
3.58
a. K+; it has a noble gas electron configuration
b. O2–; it has a noble gas electron configuration
c. Br–; it has a noble gas electron configuration
3.60
a. [Ar] or 1s2, 2s2, 2p6, 3s2, 3p6
b. [Ar] or 1s2, 2s2, 2p6, 3s2, 3p6
3.62
a.
b.
c.
d.
3.64
From Figure 3.9, He.
3.66
a. (Smallest) Li, Na, K (Largest)
b. (Smallest) F, Br, Cl (Largest)
c. (Smallest) Se, S, O (Largest)
3.68
A negative ion is always larger than its parent atom because the positive charge
is the nucleus is shared among a larger number of electrons in the ion. Each
electron moves further from the nucleus and the volume of the ion increases.
3.70
A sodium ion has a complete octet of electrons and an electron configuration
resembling its nearest noble gas.
(Smallest) Cl, S, P, Si, Al (Largest)
(Smallest) B, Al, Ga, In, Tl ( Largest)
(Smallest) Be, Mg, Ca, Sr, Ba (Largest)
(Smallest) N, P, As, Sb, Bi (Largest)
Structure and Properties of Ionic and Covalent Compounds
4.1
Chemical Bonding
Lewis Symbols
Principal Types of Chemical Bonds: Ionic and Covalent
Polar Covalent Bonding and Electronegativity
4.2
Naming Compounds and Writing Formulas of Compounds
Ionic Compounds
Covalent Compounds
A HUMAN PERSPECTIVE: Origin of the Elements
4.3
Properties of Ionic and Covalent Compounds
Physical State
Melting and Boiling Points
Structure of Compounds in the Solid State
Solutions of Ionic and Covalent Compounds
4.4
Drawing Lewis Structures of Molecules and Polyatomic Ions
Lewis Structures of Molecules
A CLINICAL PERSPECTIVE: Blood Pressure and the Sodium Ion / Potassium Ion Ratio
Lewis Structures of Polyatomic Ions
Lewis Structure, Stability, Multiple Bonds, and Bond Energies
Lewis Structures and Resonance
Lewis Structures and Exceptions to the Octet Rule
Lewis Structures and Molecular Geometry; VSEPR Theory
Lewis Structures and Polarity
4.5
Properties Based on Electronic Structure and Molecular Geometry
Solubility
Boiling Points of Liquids and Melting Points of Solids
Instructional Objectives
Conceptual Objectives
• Describe the essential differences between ionic, covalent, and polar covalent compounds.
• Recognize the differences in physical state, melting and boiling points, solid-state structure, and
solution chemistry that result from differences in bonding.
• Know the relationship between stability and bond energy.
• Understand the role that molecular geometry plays in determining the solubility and melting and
boiling points of compounds.
Performance Objectives
• Write the formulas of compounds when provided with the name of the compound.
• Name common inorganic compounds using systematic names, and recognize the common names
of frequently used substances.
• Draw Lewis structures for covalent compounds and complex inorganic ions.
• Predict the geometry of molecules and ions using the octet rule and Lewis structure.
• Use the principles of VSEPR theory and molecular geometry to predict relative melting points,
boiling points, and solubilities of compounds.
Health Applications
• Describe the way in which the sodium/potassium ratio affects blood pressure.
• Provide examples of low-sodium, high-potassium foods and high-sodium, low-potassium foods.
In-Chapter Examples
Example 4.1:
Example 4.2:
Example 4.3:
Example 4.4:
Example 4.5:
Example 4.6:
Example 4.7:
Example 4.8:
Example 4.9:
Example 4.10
Example 4.11:
Example 4.12:
Example 4.13:
Example 4.14:
Example 4.15:
Predicting the formula of an ionic compound (sodium oxide).
Predicting the formula of an ionic compound (aluminum oxide).
Writing a formula when given the name of the compound (sodium sulfate).
Writing a formula when given the name of the compound (ammonium sulfide).
Naming a covalent compound (dinitrogen tetroxide).
Writing the formula of a covalent compound (nitrogen monoxide).
Writing the formula of a covalent compound (dinitrogen tetroxide).
Drawing Lewis structures of covalent compounds (carbon dioxide) .
Drawing Lewis structures of covalent compounds (ammonia).
Drawing Lewis structures of polyatomic cations (the ammonium ion).
Drawing Lewis structures of polyatomic anions (the sulfate ion).
Drawing Lewis structures of polyatomic anions (the acetate ion).
Drawing resonance hybrids of covalently bonded compounds (the nitrate ion).
Drawing Lewis structures of covalently bonded compounds that are exceptions to
the octet rule (beryllium hydride).
Drawing Lewis structures of covalently bonded compounds that are exceptions to
the octet rule (phosphorous pentafluoride).
Chapter Overview
Chemical Bonding
When two atoms are joined to make a chemical compound, the force of attraction between the two
atoms is the chemical bond. In ionic bonding, electrons are transferred before bond formation, forming an
ion pair. In covalent bonding, electrons are shared between atoms in the bonding process. Polar covalent
bonding, like covalent bonding, is based on the concept of electron sharing; however, the sharing is
unequal and based on the electronegativity difference between joined atoms. The Lewis symbol, showing
only valence electrons, is a convenient way of representing atoms singly or in combination.
Naming Compounds and Writing Formulas of Compounds
The "shorthand" symbol for a compound is its formula. The formula identifies the number and type
of atoms in the compound. Ions are classified as monatomic or polyatomic: names of ionic compounds
are derived from the names of their ions. In the Stock system a Roman numeral indicates the charge of
the cation. This system is preferred over the older "common nomenclature" system.
Most covalent compounds are formed by the reaction of non-metals. Covalent compounds exist as
molecules and are named using prefixes that denote the number of each element present in the compound.
Many familiar compounds have common names. It is useful to correlate both systematic and
common names with the corresponding molecular formula.
Properties of Ionic and Covalent Compounds
In the solid state, covalently bonded molecules are discrete units and have less tendency to form an
extended structure, while ionic compounds form a crystal lattice. The melting and boiling temperatures
for ionic compounds are generally higher than those of covalent compounds. Ionic solids are crystalline,
whereas covalent solids may be either crystalline or amorphous. Many ionic solids dissolve in water,
dissociating into positive and negative ions (an electrolytic solution). Because these ions can carry
(conduct) a current of electricity, they are called electrolytes. Covalent solids in solution usually retain
their neutral character and are nonelectrolytes.
Drawing Lewis Structures of Molecules and Polyatomic Ions
The electronic structure of atoms, ions, and molecules, which is closely allied to the properties of
these substances, can conveniently be represented using Lewis structures, or electron-dot diagrams based
on the octet rule.
Concepts of stability and polarity are better understood when viewed from the perspective of a Lewis
structure. The stability of a covalent compound is related to the bond energy. The magnitude of the bond
energy increases and the bond length decreases in the order: single bond > double bond > triple bond.
The valence shell electron pair repulsion (VSEPR) theory helps to explain (or predict) molecular
geometry, including linear, trigonal planar, and tetrahedral arrangements of attached atoms. Exceptions to
the octet rule exist.
A polar covalent molecule has at least one polar covalent bond. An understanding of the concept of
electronegativity helps to assess the polarity of a bond. A molecule containing only nonpolar bonds must
be nonpolar. a molecule containing polar bonds may be polar or nonpolar, depending on the relative
position of the bonds.
Properties Based on Electronic Structure and Molecular Geometry.
Attractions between molecules are intermolecular forces. Intramolecular forces are the attractive
forces within molecules. It is the intermolecular forces which determine such properties as the solubility
of one substance in another and the freezing and boiling points of liquids.
Solubility is the maximum amount of solute that dissolves in a given amount of solvent at a specified
temperature. The rule of "like dissolves like" describes solubility of covalent compounds.
The amount of energy required to boil a liquid depends upon the strength of the intermolecular
attractive forces in the liquid, which, in turn, depends on the polarity of the molecules. As a general rule,
polar compounds have strong intermolecular forces, and their boiling and melting points tend to be higher
than nonpolar compounds of similar molecular mass.
Hints for Faster Coverage
Sections 4.1- 4.3 provide an overview of ionic and covalent compounds. In survey coverage, section
4.4 may be briefly mentioned, without describing molecular geometry in terms of Lewis structures.
Section 4.5 may be delayed and discussed in conjunction with solution chemistry in Chapter Seven.
Suggested Problem Sets
Chemical Bonding - 29, 31, 33
Naming Compounds and Writing Formulas of Compounds - 35, 37, 39, 41, 43, 45, 47
Properties of Ionic and Covalent Compounds - 49, 51
Drawing Lewis Structures of Ions and Molecules- 53, 55, 57, 59, 61, 63, 65
Properties Based on Electronic Structure and Molecular Geometry - 67, 69
In-Chapter Perspectives
A HUMAN PERSPECTIVE: Origin of the Elements. This perspective describes the essential features of
the "big bang" theory. Students often have little idea of the origin of elements and many believe
chemicals are substances only found in chemical laboratories. this perspective may help to spark a
real-world connection.
A CLINICAL PERSPECTIVE: Blood Pressure and the Sodium Ion/Potassium Ion Ratio. This
perspective discusses the relationship between dietary sodium intake, its relationship to potassium levels,
and, ultimately, to high blood pressure.
Additional Perspectives
Useful additional discussions might expand the relationship between structure and properties.
Examples could relate to a variety of real-life situations. Environmental considerations include the
selective movement of waste (based on polarity and solubility) from landfills to groundwater to drinking
water supplies. Biochemical examples include fat-soluble vs. water-soluble vitamins. The area of
selective catalysis affords an excellent opportunity to demonstrate structure/reactivity relationships.
Critical Thinking Problems
1.
This problem can be used to initiate a discussion of the tremendous impact that water has on our
environment; this effect is a direct consequence of the structure of the water molecule.
2.
The Lewis diagram approach is a useful tool for predicting stable combinations of atoms. On the
other hand, many exceptions exist. This problem can be used to expand this concept.
3.
The relationship of resonance and stability is emphasized.
4.
This is a useful exercise for the student to predict a structure and geometry for a molecule and
subsequently use this model to predict properties.
Chapter 4
Structure and Properties of
Ionic and Covalent Compounds
Solutions to the Even-Numbered Problems
In-Chapter Questions and Problems
4.2
a.
b.
c.
KCl (one K+ and one Cl-)
MgBr2 (one Mg2+ and two Br-)
Mg3N2 (three Mg2+ and two N3-)
4.4
a.
b.
c.
lithium carbonate
iron(II) bromide or iron dibromide
copper(II) sulfate or copper sulfate
4.6
a.
b.
c.
Na3PO4
KBr
Fe(NO3)2
4.8
a.
b.
c.
d.
hydrogen sulfide
carbon disulfide
phosphorus pentachloride
diphosphorus pentoxide
4.10
a.
b.
NF3
CO
4.12
a.
H H
H
C C H
H H
4.14
b.
:N:::N:
a.
CN–
[:C:::N:]
[Note: 5 valence electrons from nitrogen, 4 from carbon, and one from the -1 charge equals
10 valence electrons total.]
b.
CO32–
2-
O
O C O
[Note: 4 valence electrons from carbon, 6 from each of three oxygen atoms, and two for the
-2 charge equal 24 valence electrons total.]
4.16
a.
The hydrogen atom is less electronegative than the sulfur atom, so a likely skeletal structure
is:
H–S
The total number of valence electrons in the hydrogen sulfide ion is:
1 hydrogen atom x 1 valence electron
1 sulfur atom x 6 valence electrons
and 1 negative charge
= 1 electron
= 6 electrons
= 1 electron
for a total of 8 valence electrons.
These electrons can be distributed to produce the following Lewis structure:
H S
b.
The peroxide ion must have the following skeletal structure:
O–O
The total number of valence electrons in the peroxide ion is:
2 oxygen atoms x 6 valence electrons
and 2 negative charges
= 12 electrons
= 2 electrons
for a total of 14 valence electrons.
These may be distributed to produce the following structure:
2-
O O
4.18
Look at the Lewis structure for
a.
N2
:N:::N:
b.
Cl2
- a triple bond between the nitrogen atoms.
Cl Cl
- a single bond between the chlorine atoms.
The bond strength parallels the bond order. N2 has a bond order of 3, and is a stronger bond than
that found in Cl2, a single bond, bond order of 1.
4.20
Since both S and Se are in the same family (Group VIA), they have the same number of valence
electrons, and therefore, should form the same kinds of bonds.
4.22
a.
C2H4
H
C
H
H
H
H
C C
C
H
H
H
Three groups around each carbon; trigonal planar around each carbon.
b.
H:C:::C:H
H–CC–H
Two groups around each carbon; linear geometry
4.24
4.26
a.
Si – Cl
Chlorine is more electronegative than silicon; the bond is polar and the
electrons are pulled toward chlorine.
b.
S – Cl
Chlorine is more electronegative than sulfur; the bond is polar and the
electrons are pulled toward chlorine.
c.
H–C
Carbon is more electronegative than hydrogen; the bond is polar and the
electrons are pulled toward carbon.
d.
C–C
There is no electronegativity difference between two identical atoms; the bond
is nonpolar.
a.
CO2
O
C
O
There are two identical groups around the central atom. The molecule has linear geometry
and the polarity of the groups cancel. The molecule is nonpolar.
b.
SCl2
Cl
S
Cl
Sulfur and oxygen are in the same group of the periodic table. SCl 2 is analogous to H2O: a
bent molecule. The effect of two lone pairs causes the molecule to be polar.
c.
Br-Cl
Br Cl
The Br-Cl bond is polar due to the electronegativity difference between bromine and
chlorine. Since Br-Cl is the only bond in the molecule, the molecule is polar.
d.
CS2
S
C
S
There are two identical groups around the central atom. The molecule has linear geometry
and the polarity of the groups cancel. The molecule is nonpolar.
4.28
a.
C2H6 is nonpolar, heavier molecule
CH4 is nonpolar, lighter molecule
The heavier nonpolar molecule, C2H6, is predicted to have the higher melting and boiling
point.
b.
CO is polar
NO is polar
CO is the more polar molecule because there is a greater difference in the electronegativities
of carbon and oxygen than that which exists between nitrogen and oxygen. Therefore CO is
expected to have the higher melting and boiling points.
c.
F2 is nonpolar, lighter molecule
Br2 is nonpolar, heavier molecule
The heavier nonpolar molecule, Br2, is predicted to have the higher melting and boiling
point.
d.
CHCl3 is polar.
It has a higher melting point and boiling point when compared to CCl4 which is nonpolar.
End-of-Chapter Questions and Problems
4.30
a.
b.
c.
d.
Ionic (NaCl is formed from a metal and a nonmetal.)
Covalent (CO is formed from two nonmetals.)
Covalent (ICl is formed from two nonmetals.)
Covalent (H2) is formed from two identical atoms.)
4.32
a.
2-
2 Na +
+
O
2 Na
+
O
b.
2-
2 Na +
4.34
a.
2 Na+ +
S
HNO3
H
N
Ox3=6x3
S
= 1 valence electron
= 5 valence electrons
= 18 valence electrons
Total of 24 valence electrons
O
N O H
This structure satisfies the octet rule for N and O.
O
b.
CCl4
C
= 4 valence electrons
4 x Cl = 4 x 7
= 28 valence electrons
Total of 32 valence electrons
Cl
Cl C Cl
This structure satisfies the octet rule for C and Cl.
Cl
c.
PBr3
P
= 5 valence electrons
3 x Br = 3 x 7
= 21 valence electrons
Total of 26 valence electrons
Br
P
Br
This structure satisfies the octet rule for P and Br.
Br
4.36
a.
b.
c.
d.
e.
sulfide ion
chloride ion
carbonate ion
ammonium ion
acetate ion
4.38
4.40
4.42
4.44
4.46
4.48
a.
SO42–
c.
PO43–
b.
NO3–
d.
HCO3–
a.
AgCN
[one 1+ ion cancels one 1- ion]
b.
NH4Cl
[one 1+ ion cancels one 1- ion]
c.
Ag2O
[two 1+ ions cancel one 2- ion]
d.
MgCO3
[one 2+ ion cancels one 2- ion]
e.
Mg(HCO3)2 [one 2+ ion cancels two 1- ions]
a.
Nitrogen dioxide
b.
Sulfur trioxide
c.
Phosphorus trichloride
d.
Dinitrogen tetraoxide
e.
Carbon tetrachloride
a.
CO or CO2
c.
CaO
b.
SO2 or SO3
d.
SiH4
a.
NH4I
c.
NH4C2H3O2
b.
(NH4)2SO4
d.
NH4CN
a.
sodium hypochlorite
c.
sodium chlorate
b.
sodium chlorite
d.
sodium perchlorate
4.50
The attractive forces among positive and negative ions are quite strong; a great deal of energy is
needed to overcome these forces. As a result, the melting points of ionic compounds are quite
high when compared to those of covalent compounds that have weaker attractive forces among
the uncharged molecules.
4.52
Water will have a higher boiling point. Water is a polar molecule with strong intermolecular
attractive forces, whereas carbon tetrachloride is a nonpolar molecule with weak intermolecular
attractive forces. More energy, hence, a higher temperature is required to overcome the attractive
forces among the water molecules.
4.54
a.
Be
b.
B
F
c.
S
d.
2-
O
4.56
a.
Be2+
b.
Al3+
c.
d.
2-
S
4.58
Resonance is important in bonding because it accounts for the unusual stability and bond lengths
of some molecules. For instance, for the nitrite ion, two resonance forms can be drawn:
O N O
and
O
N O
If resonance did not occur, the bond lengths between N and O would not be identical. The
nitrogen oxygen double bond length would be shorter than the nitrogen oxygen single bond
length. Experimentally, we find the nitrogen oxygen bond lengths to be identical. This is
attributed to the resonance hybrid.
O
4.60
H C H
formaldehyde
H H
H
H
C C N
H
H H
4.62
ethylamine
4.64
a.
Cl and Cl
nonpolar; since there is no electronegativity difference between the atoms
sharing the electrons, the electrons are shared equally.
b.
H and H
nonpolar; there is no electronegativity difference between the atoms
sharing the electrons, the electrons are shared equally.
c.
C and H
nonpolar or only very slightly polar; the electronegativity difference
between carbon and hydrogen is negligibly small, thus the electrons are
shared essentially equally.
d.
Li and F
ionic; the bond is between a metal and a nonmetal.
e.
O and O
nonpolar; since there is no electronegativity difference between the atoms
sharing the electrons, the electrons are shared equally.
a.
Cl and Cl
4.66
b.
H and H
c.
C and H
H
H C H
Cl Cl
H
H
d.
Li and F form an ionic compound, lithium fluoride.
e.
O and O
O
O
H
4.68
The rule of "like dissolves like" applies to the question of solubility. Molecules of similar
polarity will be mutually soluble. Water is a polar molecule; therefore one would predict that
polar molecules would dissolve in water.
4.70
Polar compounds have strong intermolecular attractive forces. High temperatures are needed to
overcome these forces and convert the liquid to a gas; hence, we predict higher boiling points for
polar compounds when compared to non-polar compounds.
1. Which two scientists in 1869 arranged the elements in order of
increasing atomic masses to form a precursor of the modern periodic
table of elements?
2. Who stated that the elements, when arranged according to their
atomic
masses, showed a distinct periodicity of their properties?
3. In the modern periodic table, the elements are arranged according
to
what system?
4. The modern periodic law states that the physical and chemical
properties of the elements are periodic functions of what property?
5. What do we call the horizontal row of elements on the periodic
table?
6. How many periods are found on the periodic table?
7. Which period contains the element sodium?
8. What do we call the columns of elements on the periodic table?
9. What number for an atom gives the number of electrons and protons
found
in that atom?
10. Where are the alkaline earth metals located on the periodic table?
11. What is the general name given to the elements of Group VIIA (17)?
12. What term is used for the elements straddling the "staircase"
boundary
between the metals and nonmetals?
13. For a representative element, how can we deduce the number of
valence
electrons in a neutral atom from the position of the element in the
Periodic Table?
14. How many orbitals are in an s sublevel? How many in a p sublevel?
15. In what way(s) are the three orbitals in the 2p sublevel similar;
in
what way(s) are they different?
16. What requirement must be met in order for two electrons to coexist
in
the same orbital?
17. State the Aufbau Principle.
18. How many electrons are present in an atom of silicon?
19. Give the electronic configuration in an atom of argon, element
number
18.
20. Give the electronic arrangement in an atom of strontium, element
number
38.
21. How many electrons are present in a chloride ion?
22. State the Octet Rule.
23. Give the name of a Group IA (1) ion that has the following
electronic
arrangement: 1s22s22p6
Page 1
24. Give the name of a VIIA (17) ion that has the following electronic
arrangement: 1s22s22p63s23p6
25. What ion carries a 2- charge and is isoelectronic with K+?
26. Give the complete electronic arrangement of a sulfide ion, S2-.
27. Atoms with the biggest radii occur in the _______ _______ region
of the
Periodic Table.
28. How would you expect an Al3+ ion to compare in size with an Al
atom?
Explain why.
29. Which group of elements has the highest ionization energies? Which
group has the lowest?
30. Explain what is meant by electron affinity.
31. In Mendeleev's table of the elements, they were arranged according
to
A. atomic number
B. mass number
C. atomic mass
D. neutron number
E. density
32. The modern periodic table is arranged according to what property?
A. atomic number
B. mass number
C. atomic mass
D. neutron number
E. density
33. What do we call a complete horizontal row of elements on the
periodic
table?
A. group
B. period
C. family
D. representative elements
E. transition elements
34. What are all the elements in the A-groups often called?
A. transition elements
B. lanthanides
C. metals
D. non-metals
E. representative elements
35. Which of the following elements is a metalloid?
A. C B. Ge C. Pb D. N E. P
Page 2
36. Where are the alkali metals located on the periodic table?
A. representative elements
B. transition metals
C. Group IA (1)
D. Group IIA (2)
E. Group IIIA (3)
37. How many valence electrons are in an atom of carbon?
A. 8 B. 6 C. 4 D. 1 E. 0
38. What is the lowest energy sublevel of a principal level?
A. d B. e C. f D. s E. p
39. How many sublevels are there in the third principal energy level?
A. 3 B. 2 C. 1 D. 0 E. 4
40. How many orbitals are there in a p sublevel?
A. 2 B. 3 C. 1 D. 0 E. 4
41. Which of the following correctly gives the electron capacity of a
principal energy level in terms of the number n?
A. n B. 2n C. 2n + 2 D. n2 E. 2n2
42. What is the electron configuration of sulfur, atomic number 16?
A. 1s21p62s22p6
B. 1s22s22p62d6
C. 1s22s22p63s23p4
D. 1s22s22p63s23d4
E. 1s22s22p63s22d4
43. Which one of the following electron configurations is appropriate
for a
normal atom?
A. 1s12s1
B. 1s22s1
C. 1s22s22p8
D. 1s22s22p43s1
E. 1s22s22p63d1
44. Which of the following elements is most likely to form a 3+ ion?
A. Li B. K C. Al D. N E. Cu
Page 3
45. Give the complete electronic configuration of a sodium ion.
A. 1s22s22p5
B. 1s22s22p6
C. 1s22s22p63s1
D. 1s22s22p63s2
E. 1s22s22p63s23p64s1
46. Which of the following ions does not follow the octet rule?
A. Na+ B. Ca2+ C. Al3+ D. N3- E. Cl247. Which of the following atoms has the biggest size (radius)?
A. Na B. Al C. Cl D. Rb E. I
48. Which of the following elements has the highest ionization energy?
A. Li B. B C. O D. F E. Ne
49. Which of the following elements has the lowest ionization energy?
A. Li B. B C. O D. F E. Ne
50. The electron affinity is
A. the energy required to remove an electron from an isolated atom
B. the force between two electrons in the same orbital
C. the force between two ions of opposite charge
D. the energy released when an isolated atom gains an electron
E. the attraction of an atom for an electron in a chemical bond
51. Which one of the following elements has the highest electron
affinity?
A. Li B. K C. Kr D. O E. Cl
52. T F In Mendeleev's table, the elements were arranged according to
their atomic numbers.
53. T F There are nine periods on the periodic table.
54. T F Sulfur (S) is one of the representative elements.
55. T F Platinum (Pt) is a lanthanide element.
56. T F Tin (Sn) is a metalloid.
57. T F Valence electrons are involved when atoms form bonds.
58. T F There are a maximum of 50 electrons in principal energy level
number five.
59. T F Atoms of the noble gas elements, Group VIII A (18), do not
form
bonds with any other elements.
60. T F There are eight valence electrons in a chloride ion.
61. T F The ions formed from Group IIA (2) atoms have charges of 2+.
Page 4
62. T F Cations tend to be formed from metal atoms, while anions are
formed from non-metal atoms.
63. T F The atoms of smallest radius are those of elements in top left
hand part of the periodic table.
64. T F The halogens (Group VII A (17)) have the lowest ionization
energies of any group in the periodic table.
Page 5
Answer Key for Test "chapter3.tst", 8/17/04
No. in
Q-Bank
No. on
Test Correct Answer
3 1 1 Mendeleev and Meyer
3 2 2 Dimitri Mendeleev
3 3 3 increasing atomic number
3 4 4 atomic number
3 5 5 periods
3 6 6 seven
3 7 7 three
3 8 8 groups
3 9 9 atomic number
3 10 10 Group IIA (2)
3 11 11 halogens
3 12 12 metalloids
3 13 13 the group number is also the number of valence electrons
3 14 14 1;3
3 15 15 they have the same shape and the same energy; they are
oriented
differently in space
3 16 16 they must have opposite spins
3 17 17 Electrons occupy the available orbital of lowest energy first.
3 18 18 Fourteen
3 19 19 1s22s22p63s23p6
3 20 20 1s22s22p63s23p64s23d104p65s2
3 21 21 Eighteen
3 22 22 Elements tend to react in such a way as to attain the electron
configuration of the atoms of the noble gas nearest to them in
the Periodic Table.
3 23 23 sodium ion
3 24 24 chloride
3 25 25 S23 26 26 1s22s22p63s23p6
3 27 27 bottom left
3 28 28 The ion will be much smaller. In forming the ion, the atom
loses
all its outermost electrons. The net positive charge on the ion
ensures that all the electrons in the ion are strongly attracted
to the nucleus, keeping the ion small.
3 29 29 Group VIIIA (18) are highest; Group IA (1) are the lowest.
3 30 30 It is the energy released when a neutral atom gains an
electron
to form an anion.
3 31 31 C
3 32 32 A
3 33 33 B
3 34 34 E
3 35 35 B
3 36 36 C
3 37 37 C
3 38 38 D
3 39 39 A
3 40 40 B
3 41 41 E
3 42 42 C
3 43 43 B
3 44 44 C
3 45 45 B
3 46 46 E
Page 1
Answer Key for Test "chapter3.tst", 8/17/04
No. in
Q-Bank
No. on
Test Correct Answer
3 47 47 D
3 48 48 E
3 49 49 A
3 50 50 D
3 51 51 E
3 52 52 F
3 53 53 F
3 54 54 T
3 55 55 F
3 56 56 F
3 57 57 T
3 58 58 T
3 59 59 F
3 60 60 T
3 61 61 T
3 62 62 T
3 63 63 F
3 64 64 F
Page 2
1. In a Lewis structure, what do the dots represent?
2. Draw the Lewis structure of the bromine atom.
3. How many dots are shown in the Lewis structure for the sulfur atom?
4. What are the two principal types of bonding called?
5. Name the two classes of element which are most likely to form an
ionic
compound if they are allowed to react with each other.
6. Draw the Lewis structure of the Pb2+ ion.
7. What constitutes a covalent bond between two atoms?
8. In what way is a polar covalent bond similar to a nonpolar covalent
bond? In what way are they different?
9. What does it mean if an atom is said to have a high
electronegativity?
10. The elements with the lowest electronegativities are found in the
_______ _______ region of the periodic table.
11. Who first assigned electronegativity values to many of the
elements?
12. What do we call the three-dimensional arrangement of positive and
negative ions in an ionic solid?
13. Predict the formula of the compound formed when ions of sodium and
sulfur combine.
14. Predict the formula of the compound formed when ions of barium and
nitrogen combine.
15. What is the name of Fe2+ in the Stock system?
16. What does the suffix "-ous" on the common names of ions mean?
17. What is the term used for ions that are composed of two or more
atoms
bonded together?
18. What is the formula of the sulfate ion?
19. What is the name of the ion HCO3-?
20. What is the name of the ion NH4+?
21. Provide the name of Na3PO4.
22. What is the name of Cu2O in the Stock system?
23. Write the formula of sodium carbonate.
24. What kind of compound results when two or more different nonmetals
share electrons?
25. What kind of bonding exists in substances which consist of
discrete
molecules?
Page 1
26. Provide the formula of sulfur trioxide.
27. Write the formula of ammonia.
28. Provide the name of CCl4.
29. Provide the name of the compound whose formula is N2O5.
30. At what temperature is a liquid converted into a gas?
31. What is the term that describes a solid with no regular structure?
32. What is the term that describes a compound that, when dissolved in
water conducts an electric current?
33. What is the term that describes a compound that, when dissolved in
water does not conduct an electric current?
34. What kind of bonding is present in substances which are
nonelectrolytes?
35. How many bonding electrons are shown in the Lewis structure for
the
bicarbonate ion, HCO3-?
36. Draw the Lewis structure of methylamine, CH3NH2.
37. Draw the Lewis structure of hydrogen sulfide, H2S.
38. What is wrong with the Lewis structure shown below for sulfur
trioxide,
SO3?
SO
O
O
39. Ozone, O3, has two resonance forms. Draw them, given the skeletal
arrangement O-O-O
40. What is defined as the amount of energy needed to break a bond
holding
two atoms together?
41. What is defined as the distance of separation of two nuclei in a
covalent bond?
42. What do the letters VSEPR stand for?
43. If the shape of a molecule is trigonal planar, what are the values
of
the bond angles?
44. In the molecule AX2, the central atom A has two lone pairs of
electrons
in addition to the two bond pairs in the A-X bonds. What is the shape
of this molecule?
45. The ammonia molecule, NH3, is polar. Why does this fact suggest
that its
shape is trigonal pyramidal, rather than trigonal planar?
Page 2
46. Which of the following Lewis structures of neutral atoms is
correct?
K Sn O Ba Al
A B C D E
A. A B. B C. C D. D E. E
47. Which of the following Lewis structures of ions is incorrect?
Na Ca Sn N I
2+ 3- - + 2+
A B C D E
A. A B. B C. C D. D E. E
48. Which of the following has the greatest electronegativity?
A. H B. Cl C. O D. F E. Na
49. Which of the following has the greatest electronegativity?
A. Si B. P C. Cl D. Ar E. Br
50. In the compound CH3Cl the bond between carbon and chlorine is
A. intermolecular B. ionic
C. nonpolar covalent D. polar covalent
51. Which one of the following is NOT true about elements that form
cations?
A. The atoms lose electrons in forming ions.
B. The elements are metals.
C. They are located to the left of the periodic table.
D. They have low ionization energies.
E. They have high electron affinities.
52. Assuming reactions between the following pairs of elements, which
pair
is most likely to form an ionic compound?
A. copper and tin
B. chlorine and oxygen
C. cesium and iodine
D. carbon and chlorine
E. fluorine and iodine
53. What kind of bond results when electron transfer occurs between
atoms
of two different elements?
A. ionic
B. covalent
C. nonpolar
D. single
E. double
Page 3
54. What is the old name of Cu+?
A. cupric ion
B. cuprous ion
C. copper(I) ion
D. copper(II) ion
E. ferrous ion
55. Give the name of FeSO4 in the Stock system.
A. iron monosulfuric acid
B. iron(II) sulfate
C. iron(III) sulfate
D. ferrous sulfate
E. ferric sulfate
56. Assuming reactions between the following pairs of elements, which
pair
is most likely to form a covalent compound?
A. lithium and iodine
B. sodium and oxygen
C. calcium and chlorine
D. copper and tin
E. carbon and oxygen
57. A double bond between two atoms, A and B
A. is longer than a single bond between the same two atoms
B. has a lower bond energy than a single bond between the same two
atoms
C. arises when two electrons are transferred from A to B
D. consists of two electrons shared between A and B
E. consists of four electrons shared between A and B
58. What is the correct formula of phosphorus pentachloride?
A. PCl B. PCl3 C. PCl5 D. P2Cl5 E. P5Cl
59. What term describes the temperature at which a solid is converted
into
a liquid?
A. critical point
B. flash point
C. sublimation point
D. melting point
E. boiling point
60. What term describes a solution of a compound in water that
conducts an
electric current?
A. amorphous solution
B. an electrolyte solution
C. a nonelectrolyte solution
D. superconducting solution
E. isoelectric solution
Page 4
61. How many bonding electrons are in CO2?
A. 1 B. 2 C. 3 D. 4 E. 8
62. How many nonbonding electrons are in CH4?
A. 0 B. 1 C. 2 D. 3 E. 8
63. How many valence electrons are in SO42-?
A. 2 B. 64 C. 32 D. 12 E. 16
64. According to VSEPR theory, if the valence electrons on a central
atom
are 3 bond pairs and one nonbonding (lone) pair, the geometry (shape)
at this atom will be
A. linear
B. bent (angular)
C. trigonal planar
D. trigonal pyramidal
E. tetrahedral
65. T F In Lewis structures, the chemical symbol of an element
represents
both the nucleus and the lower energy (nonvalence) electrons.
66. T F The name of SnO2 is tin(I) oxide.
67. T F The old name of iron(III) chloride is ferrous chloride.
68. T F The are three atoms of iodine represented in the formula
NaIO3.
69. T F In solid NaCl, no molecules of NaCl exist.
70. T F Ionic solids are amorphous.
71. T F Molecular compounds usually involve ionic bonding.
72. T F As a rule, ionic compounds tend to have lower melting and
boiling
points than covalent compounds consisting of small molecules.
73. T F In the water molecule, the oxygen atom is an exception to the
octet rule.
74. T F The NO2 molecule can never satisfy the octet rule.
75. T F Six electrons shared between two atoms corresponds to a bond
order of three.
76. T F Resonance occurs when two or more different, valid Lewis
structures can be drawn for a molecule+.
77. T F The existence of resonance makes a molecule less stable than
would otherwise be the case.
78. T F Because the C-H bond in methane is polar, the CH4 molecule
will
also be polar.
79. T F Chemical bonds are intramolecular forces.
Page 5
80. T F In determining properties such as solubility, melting point
and
boiling point, intramolecular forces are more important than
intermolecular forces.
81. T F As a rule, a polar substance will be a good solvent for
nonpolar
solutes, and vice versa.
Page 6
Answer Key for Test "chapter4.tst", 8/17/04
No. in
Q-Bank
No. on
Test Correct Answer
4 1 1 valence electrons
4 2 2
Br
4 3 3 six
4 4 4 ionic bonding and covalent bonding
4 5 5 metal, nonmetal
4 6 6 Pb 2
4 7 7 a shared pair of electrons
4 8 8
In each case, the bond consists of an electron pair shared between the
bonded atoms. The difference is that the sharing is unequal in the
case
of the polar covalent bond, equal in a nonpolar covalent bond.
4 9 9 The atom has a strong attraction for shared electron pairs
(electrons in covalent bonds).
4 10 10 bottom left
4 11 11 Pauling
4 12 12 crystal lattice
4 13 13 Na2S
4 14 14 Ba3N2
4 15 15 iron(II) ion
4 16 16 lower positive charge
4 17 17 polyatomic
4 18 18 SO424 19 19 hydrogen carbonate
4 20 20 ammonium
4 21 21 sodium phosphate
4 22 22 copper(I) oxide
4 23 23 Na2CO3
4 24 24 covalent
4 25 25 Covalent
4 26 26 SO3
4 27 27 NH3
4 28 28 carbon tetrachloride
4 29 29 dinitrogen pentoxide
4 30 30 boiling point
4 31 31 amorphous
4 32 32 electrolyte
4 33 33 nonelectrolyte
4 34 34 covalent
4 35 35 ten
4 36 36
CN
HH
H
HH
4 37 37 H S H
4 38 38
Page 1
Answer Key for Test "chapter4.tst", 8/17/04
No. in
Q-Bank
No. on
Test Correct Answer
The structure shows 26 valence electrons, but there should only be 24.
4 39 39
OOOOOO
4 40 40 bond energy
4 41 41 bond length
4 42 42 Valence Shell Electron Pair Repulsion
4 43 43 120°
4 44 44 bent or angular
4 45 45 If the molecule were trigonal planar, the symmetry would
result
in a nonpolar molecule. The centers of positive and negative
charge would coincide.
4 46 46 C
4 47 47 C
4 48 48 D
4 49 49 C
4 50 50 D
4 51 51 E
4 52 52 C
4 53 53 A
4 54 54 B
4 55 55 B
4 56 56 E
4 57 57 E
4 58 58 C
4 59 59 D
4 60 60 B
4 61 61 E
4 62 62 A
4 63 63 C
4 64 64 D
4 65 65 T
4 66 66 F
4 67 67 F
4 68 68 F
4 69 69 T
4 70 70 F
4 71 71 F
4 72 72 F
4 73 73 F
4 74 74 T
4 75 75 T
4 76 76 T
4 77 77 F
4 78 78 F
4 79 79 T
4 80 80 F
4 81 81 F
Page 2