Solution Manual for Chemistry, The Central Science 13th Edition by Theodore E. Brown , H. Eugene LeMay, Bruce E. Bursten

Solution Manual for Chemistry, The Central Science 13th Edition by Theodore E. Brown , H. Eugene LeMay, Bruce E. Bursten

Chapter 7. Periodic Properties of the Elements
Media Resources
Important Figures and Tables: Section:
Figure 7.3 Effective Nuclear Charge 7.2 Effective Nuclear Charge
Figure 7.4 Comparison of 1s, 2s, and 2p Radial 7.3 Sizes of Atoms and Ions
Probability Functions
Figure 7.5 Variations in Effective Nuclear Charge 7.3 Sizes of Atoms and Ions
for Period 2 and Period 3 Elements
Figure 7.7 Trends in Bonding Atomic Radii for 7.3 Sizes of Atoms and Ions
Periods 1 through 5
Figure 7.8 Cation and Anion Size 7.3 Sizes of Atoms and Ions
Table 7.2 Successive Values of Ionization Energies, 7.4 Ionization Energy
I, for the Elements Sodium Through Argon
(kJ/mol)
Figure 7.10 Trends in First Ionization Energies of 7.4 Ionization Energy
the Elements
Figure 7.12 Electron Affinity in kJ/mol for Selected 7.5 Electron Affinity
s- and p-Block Elements
Figure 7.13 Metals, Metalloids, and Nonmetals 7.6 Metals, Nonmetals, and Metalloids
Figure 7.15 Representative Oxidation States of the 7.6 Metals, Nonmetals, and Metalloids
Elements
Table 7.4 Some Properties of the Alkali Metals 7.7 Trends for Group 1A and 2A Metals
Table 7.5 Some Properties of the Alkaline Earth 7.7 Trends for Group 1A and 2A Metals
Metals
Table 7.6 Some Properties of the Group 6A Elements 7.8 Trends for Selected Nonmetals
Table 7.7 Some Properties of the Halogens 7.8 Trends for Selected Nonmetals
Table 7.8 Some Properties of the Noble Gases 7.8 Trends for Selected Nonmetals
Activities: Section:
Periodic Table 7.1 Development of the Periodic Table
Ionization Energies 7.4 Ionization Energy
Animations: Section:
Periodic Properties 7.1 Development of the Periodic Table
Effective Nuclear Charge 7.2 Effective Nuclear Charge
Periodic Trends: Atomic Radii 7.3 Sizes of Atoms and Ions
Gain and Loss of Electrons 7.3 Sizes of Atoms and Ions
Ionization Energy 7.4 Ionization Energy
Periodic Trends: Ionization Energies 7.4 Ionization Energy
Electron Affinity 7.5 Electron Affinity
Periodic Trends: Electron Affinity 7.5 Electron Affinity
Periodic Trends: Acid-Base Behavior of Oxides 7.6 Metals, Nonmetals, and Metalloids
Movies: Section:
Sodium and Potassium in Water 7.7 Trends for Group 1A and 2A Metals
Flame Tests for Metals 7.7 Trends for Group 1A and 2A Metals
Periodic Properties of the Elements 91
Physical Properties of the Halogens 7.8 Trends for Selected Nonmetals
VCL Simulations: Section:
Flame Tests for Metals 7.7 Trends for Group 1A and 2A Metals
Emission Spectra of Sodium and Mercury 7.7 Trends for Group 1A and 2A Metals
Other Resources
Further Readings: Section:
Chemical and Engineering News, September 8, 7.1 Development of the Periodic Table
2003
It’s Elemental! 7.1 Development of the Periodic Table
Using the Learning Cycle to Introduce Periodicity 7.1 Development of the Periodic Table
The Nuts and Bolts of Chemistry 7.1 Development of the Periodic Table
Mendeleev and Moseley: The Principal 7.1 Development of the Periodic Table
Discoverers of the Periodic Law
Mendeleev’s Other Predictions 7.1 Development of the Periodic Table
Atomic Numbers Before Moseley 7.1 Development of the Periodic Table
D. I. Mendeleev and the English Chemists 7.1 Development of the Periodic Table
The Evolution of the Periodic System 7.1 Development of the Periodic Table
Periodic Table of Elemental Abundance 7.1 Development of the Periodic Table
A Different Approach to a 3-D Periodic System 7.1 Development of the Periodic Table
Including Stable Isotopes
Screen Percentages Based on Slater Effective 7.2 Effective Nuclear Charge
Nuclear Charge as a Versatile Tool for
Teaching Periodic Trends
Pictorial Analogies VI: Radial and Angular Wave 7.3 Sizes of Atoms and Ions
Function Plots
Using Balls from Different Sports to Model the 7.3. Sizes of Atoms and Ions
Variation of Atomic Sizes
Periodic Contractions Among the Elements; Or, 7.3 Sizes of Atoms and Ions
On Being the Right Size
Ionization Energies of Atoms and Atomic Ions 7.4 Ionization Energy
Trends in Ionization Energy of Transition-Metal 7.4 Ionization Energy
Elements
Periodicity in the Acid-Base Behavior of Oxides 7.6 Metals, Nonmetals, and Metalloids
and Hydroxides
Metalloids 7.6 Metals, Nonmetals, and Metalloids
A Variation on the Demonstration of the Properties 7.7 Trends for Group 1A and 2A Metals
of the Alkali Metals
A Little Lithium May be Just What the Doctor 7.7 Trends for Group 1A and 2A Metals
Ordered
The Legend of Dr. Pepper/Seven-Up 7.7 Trends for Group 1A and 2A Metals
Update on Intake; Calcium Consumption Low 7.7 Trends for Group 1A and 2A Metals
Life, Death, and Calcium 7.7 Trends for Group 1A and 2A Metals
A Second Note on the Term ‘Chalcogen’ 7.8 Trends for Selected Nonmetals
Allotropes and Polymorphs 7.8 Trends for Selected Nonmetals
The Origin of the Term Allotrope 7.8 Trends for Selected Nonmetals
Aqueous Hydrogen Peroxide: Its Household Uses 7.8 Trends for Selected Nonmetals
And Concentration Units
92 Chapter 7
The Chemistry of Swimming Pool Maintenance 7.8 Trends for Selected Nonmetals
Live Demonstrations: Section:
Halogens Compete for Electrons 7.5 Electron Affinity
Acidic and Basic Properties of Oxides 7.6 Metals, Nonmetals, and Metalloids
Disappearing Ink 7.6 Metals, Nonmetals, and Metalloids
A Dramatic Flame Test Demonstration 7.7 Trends for Group 1A and 2A Metals
Simple Flame Test Techniques Using Cotton Swabs 7.7 Trends for Group 1A and 2A Metals
Producing Hydrogen Gas from Calcium Metal 7.7 Trends for Group 1A and 2A Metals
Preparation and Properties of Oxygen 7.8 Trends for Selected Nonmetals
Plastic Sulfur 7.8 Trends for Selected Nonmetals
Periodic Properties of the Elements 93
Chapter 7. Periodic Properties of the Elements
Common Student Misconceptions
• Students have difficulty with the concepts of shielding and effective nuclear charge. As you move to
the right in a period, shielding does not increase appreciably but the nuclear charge does. Therefore,
effective nuclear charge increases steadily as you move to the right along the period.
• Students are confused why, within a period, atomic radii decrease with increasing atomic number.
• Students often do not understand slight irregularities in periodic trends for elements in each row after
each ns subshell becomes filled, and after np and (n – 1)d subshells become half-filled.
• Students often have problems with the signs of electron affinities; in particular, why group 1A metals
have negative (exothermic) electronegativities.
• Students are often confused regarding the placement of hydrogen on the periodic table; despite its
common placement in column 1A, hydrogen is a nonmetal.
• Students often confuse behavior of elements in aqueous phase with periodic properties determined
in gas phase (ionization energy, electron affinity) or in solid phase (ionic radius).
• Students often confuse isoelectronic species with those with the same number of valence electrons.
Teaching Tips
• Students need to be shown how position on the periodic table and electron configurations can be used
to highlight periodic properties.
• Emphasize the periodic table as an organizational tool; it will help students recall chemical facts.
• Students find the descriptive chemistry/group trends a bit overwhelming at first.
• Live demonstrations, CD videos, and web-based animations are very helpful in stimulating student
interest in the group trends.
Lecture Outline
7.1 Development of the Periodic Table1,2,3,4,5,6,7,8,9,10,11,12
• The periodic table is the most significant tool that chemists use for organizing and recalling chemical
facts.
• Elements in the same column contain the same number of outer-shell electrons or valence electrons.
• How do we organize the different elements in a meaningful way that will allow us to make
predictions about undiscovered elements?

1
“Periodic Properties” Animation from Instructor’s Resource CD/DVD
2
“Periodic Table” Activity from Instructor’s Resource CD/DVD
3 September 8, 2003 issue of Chemical and Engineering News from Further Readings
4
“It’s Elemental!” from Further Readings
5
“The Nuts and Bolts of Chemistry” from Further Readings
6
“Mendeleev and Moseley: The Principal Discoverers of the Periodic Law” from Further Readings
7
“The Evolution of the Periodic System” from Further Readings
8
“Mendeleev’s Other Prediction” from Further Readings
9
“Atomic Number Before Moseley” from Further Readings
10 “D. I. Mendeleev and the English Chemists” from Further Readings
11 “Periodic Tables of Elemental Abundance” from Further Readings
12 “A Different Approach to a 3–D Periodic System Including Stable Isotopes” from Further Readings
94 Chapter 7
• Arrange elements to reflect the trends in chemical and physical properties.
• The periodic table arises from the periodic patterns in the electronic configurations of the elements.
• Elements in the same column contain the same number of valence electrons.
• The trends within a row or column form patterns that help us make predictions about chemical
properties and reactivity.
• In the first attempt Mendeleev and Meyer arranged the elements in order of increasing atomic weight.
• Certain elements were missing from this scheme.
• For example, in 1871 Mendeleev noted that As properly belonged underneath P and not Si, which
left a missing element underneath Si. He predicted a number of properties for this element.
• In 1886 Ge was discovered; the properties of Ge match Mendeleev’s predictions well.
• In the modern periodic table, elements are arranged in order of increasing atomic number.
FORWARD REFERENCES
• Periodic trends and chemical properties of nonmetals will be further discussed in Chapter 22
(section 22.1).
7.2 Effective Nuclear Charge13,14,15,16,17

• Effective nuclear charge (Zeff) is the charge experienced by an electron on a many-electron atom.
• The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the
inner electrons.
• The electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full
nuclear charge.
• The nuclear charge experienced by an electron depends on its distance from the nucleus and the
number of electrons in the spherical volume out to the electron in question.
• As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff)
decreases.
Zeff = Z – S
• As the distance from the nucleus increases, S increases and Zeff decreases.
• S is called the screening constant which represents the portion of the nuclear charge that is
screened from the valence electron by other electrons in the atom.
• The value of S is usually close to the number of core electrons in an atom.
• The Zeff for valence electrons varies with position in the periodic table
• Zeff increases as we move from left to right across a period
• Zeff increases slightly as we do from top to bottom in a group however the change is much less
than that seen as we move across a period
7.3 Sizes of Atoms and Ions18,19
• Consider a collection of argon atoms in the gas phase.
• When they undergo collisions, they ricochet apart because electron clouds cannot penetrate each
other to a significant extent.
• The apparent radius is determined by the closest distances separating the nuclei during such
collisions.
13 “Effective Nuclear Charge” Animation from Instructor’s Resource CD/DVD
14 Figure 7.3
15 Figure 7.4
16 “Screen Percentages Based on Slater Effective Nuclear Charge as a Versatile Tool for Teaching
Periodic Trends” from Further Readings
17 Figure 7.5
18 Figure 7.7
19 “Pictorial Analogies: VI: Radial and Angular Wave Function Plots” from Further Readings
Periodic Properties of the Elements 95
• This radius is the nonbonding radius.
• Nonbonding atomic radii are also called van der Waals radii.
• These are used in space-filling models to represent the sizes of different elements.
• Now consider a simple diatomic molecule.
• The distance between the two nuclei is called the bonding atomic radius.
• It is shorter than the nonbonding radius.
• If the two atoms which make up the molecule are the same, then half the bond distance is called
the covalent radius of the atom.
Periodic Trends in Atomic Radii20,21
• Atomic size varies consistently through the periodic table.
• As we move down a group the atoms become larger.
• As we move across a period atoms become smaller.
• There are two factors at work:
• the principal quantum number, n, and
• the effective nuclear charge, Zeff.
• As the principal quantum number increases (i.e., we move down a group), the distance of the
outermost electron from the nucleus becomes larger. Hence the atomic radius increases.
• As we move across the periodic table, the number of core electrons remains constant, however,
the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and
the outermost electrons. This attraction causes the atomic radius to decrease.
Periodic Trends in Ionic Radii22,23,24
• Ionic size is important:
• in predicting lattice energy and
• in determining the way in which ions pack in a solid.
• Just as atomic size is periodic, ionic size is also periodic.
• In general:
• Cations are smaller than their parent atoms.
• Electrons have been removed from the most spatially extended orbital.
• The effective nuclear charge has increased.
• Therefore, the cation is smaller than the parent atom.
• Anions are larger than their parent atoms.
• Electrons have been added to the most spatially extended orbital.
• This means total electron-electron repulsion has increased.
• Therefore, anions are larger than their parent atoms.
• For ions with the same charge, ionic size increases down a group.
• All the members of an isoelectronic series have the same number of electrons.
• As nuclear charge increases in an isoelectronic series the ions become smaller:
O2– > F–
> Na+
> Mg2+ > Al3+
FORWARD REFERENCES
• Sizes and charges of ions will be instrumental in determining lattice energies (Chapter 8).
• Structures of ionic solids in Chapter 12 (section 12.2).
• Atomic radii will affect relative strengths of binary acids of nonmetals from a given group, as
discussed in Chapter 16 (section 16.10).

20 “Periodic Trends: Atomic Radii” Animation from Instructor’s Resource CD/DVD
21 “Using Balls from Different Sports to Model the Variation of Atomic Sizes” from Further Readings
22 “Gain and Loss of Electrons” Animation from Instructor’s Resource CD/DVD
23 Figure 7.8
24 “Periodic Contractions Among the Elements; Or, On Being the Right Size” from Further Readings
96 Chapter 7
• Periodic properties of nonmetals in groups 4A-8A are tabulated throughout Chapter 22.
• Periodic properties for the first transition-series elements are shown in Chapter 23 (section
23.1).
7.4 Ionization Energy25,26
• The ionization energy of an atom or ion is the minimum energy required to remove an electron from
the ground state of the isolated gaseous atom or ion.
• The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous
atom:
Na(g)  Na+
(g) + e–
• The second ionization energy, I2, is the energy required to remove the second electron from a gaseous
ion:
Na+
(g)  Na2+(g) + e–
• The larger the ionization energy, the more difficult it is to remove the electron.
• There is a sharp increase in ionization energy when a core electron is removed.
Variations in Successive Ionization Energies27,28
• Ionization energies for an element increase in magnitude as successive electrons are removed.
• As each successive electron is removed, more energy is required to pull an electron away from an
increasingly more positive ion.
• A sharp increase in ionization energy occurs when an inner-shell electron is removed.
Periodic Trends in First Ionization Energies29,30,31
• Ionization energy generally increases across a period.
• As we move across a period, Zeff increases, making it more difficult to remove an electron.
• Two exceptions are removing the first p electron and removing the fourth p electron.
• The s electrons are more effective at shielding than p electrons. So, forming the s
2
p
0
configuration is more favorable.
• When a second electron is placed in a p orbital, the electron-electron repulsion increases.
When this electron is removed, the resulting s
2
p
3
configuration is more stable than the starting
s
2
p
4
configuration. Therefore, there is a decrease in ionization energy.
• Ionization energy decreases down a group.
• This means that the outermost electron is more readily removed as we go down a group.
• As the atom gets bigger, it becomes easier to remove an electron form the most spatially extended
orbital.
• Example: For the noble gases the ionization energies follow the order:
He > Ne > Ar > Kr > Xe
• The representative elements exhibit a larger range of values for I1 than transition metals.

25 “Ionization Energy” Animation from Instructor’s Resource CD/DVD
26 “Ionization Energies of Atoms and Atomic Ions” from Further Readings
27 “Ionization Energies” Activity from Instructor’s Resource CD/DVD
28 Table 7.2
29 Figure 7.10
30 “Periodic Trends: Ionization Energies” Animation from Instructor’s Resource CD/DVD
31 “Trends in Ionization Energy of Transition-Metal Elements” from Further Readings
Periodic Properties of the Elements 97
Electron Configurations of Ions
• These are derived from the electron configurations of elements with the required number of electrons
added or removed from the most accessible orbital.
• Li: [He]2s
1
becomes Li+
: [He]
• F: [He]2s2
2p
5
becomes F-
: [He]2s
2
2p
6
= [Ar]
• Transition metals tend to lose the valence shell electrons first and then as many d electrons as are
required to reach the desired charge on the ion.
• Thus electrons are removed from 4s before the 3d, etc.
• In other words, when writing electron configurations of transition metal cations, the order of
removal of electrons is not exactly opposite to the order in which subshells were occupied
when an electron configuration of the parent atom was written.
FORWARD REFERENCES
• Octet rule will be introduced in Chapter 8.
• Discussion of electron configurations of the representative elements and transition metals will
continue in Chapter 8.
• Photoionization processes and ionization energies will be linked together in Chapter 18.
(section 18.1).
7.5 Electron Affinity32,33,34,35
• Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion.
• Electron affinity and ionization energy measure the energy changes of opposite processes.
• Electron affinity: Cl(g) + e–  Cl–
(g) ∆E = –349 kJ/mol
• Ionization energy: Cl(g)  Cl+
(g ) + e–
∆E = 1251 kJ/mol
• Electron affinity can either be exothermic (as the above example) or endothermic:
Ar(g) + e–  Ar– (g) ∆E > 0
• Look at electron configurations to determine whether electron affinity is positive or negative.
• The extra electron in Ar needs to be placed in the 4s orbital which is significantly higher in
energy than the 3p orbital.
• The added electron in Cl is placed in the 3p orbital to form the stable 3p
6
electron configuration.
• Electron affinities do not change greatly as we move down in a group.
7.6 Metals, Nonmetals and Metalloids36,37
• Metallic character refers to the extent to which the element exhibits the physical and chemical
properties of metals.
• Metallic character increases down a group.
• Metallic character decreases from left to right across a period.

32 “Electron Affinity” Animation from Instructor’s Resource CD/DVD
33 “Periodic Trends: Electron Affinity” Animation from Instructor’s Resource CD/DVD
34 Figure 7.12
35 “Halogens Compete for Electrons” from Live Demonstrations
36 Figure 7.13
37 “Acidic and Basic Properties of Oxides” from Live Demonstrations
98 Chapter 7
Metals38,39,40,41,42

• Metals are shiny and lustrous, malleable and ductile.
• Metals are solids at room temperature (exception: mercury is liquid at room temperature; gallium and
cesium melt just above room temperature) and have very high melting temperatures.
• Metals tend to have low ionization energies and tend to form cations easily.
• Metals tend to be oxidized when they react.
• Compounds of metals with nonmetals tend to be ionic substances.
• Metal oxides form basic ionic solids.
• Most metal oxides are basic:
Metal oxide + water  metal hydroxide
Na2O(s) + H2O(l)  2NaOH(aq)
• Metal oxides are able to react with acids to form salts and water:
Metal oxide + acid  salt + water
NiO(s) + 2HNO3(aq)  Ni(NO3)2(aq) + H2O(l)
Nonmetals
• Nonmetals are more diverse in their behavior than metals.
• In general, nonmetals are nonlustrous, are poor conductors of heat and electricity, and exhibit lower
melting points than metals.
• Seven nonmetallic elements exist as diatomic molecules under ordinary conditions:
• H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
• When nonmetals react with metals, nonmetals tend to gain electrons:
Metal + nonmetal  salt
2Al(s) + 3Br2(l)  2AlBr3 (s)
• Compounds composed entirely of nonmetals are molecular substances.
• Most nonmetal oxides are acidic:
Nonmetal oxide + water  acid
CO2(g) + H2O(l)  H2CO3(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
• Nonmetal oxides react with bases to form salts and water:
Nonmetal oxide + base  salt + water
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
Metalloids43
• Metalloids have properties that are intermediate between those of metals and nonmetals.
• For example, Si has a metallic luster but it is brittle.
• Metalloids have found fame in the semiconductor industry.
FORWARD REFERENCES
• The role of metals and metalloids in semiconductors will be discussed in Chapter 12 (section
12.7).
• Arrhenius, Brønsted-Lowry and Lewis acids and bases will be discussed in Chapter 16.
• Acids and bases as well as reactions between them will be discussed in Chapter 16.
• An in-depth discussion of nonmetals will be provided in Chapter 22.
38 Figure 7.15
39 “Periodicity in the Acid-Base Behavior of Oxides and Hydroxides” from Further Readings
40 “Periodic Trends: Acid-Base Behavior of Oxides” Animation from Instructor’s Resource CD/DVD
41 “Acidic and Basic Properties of Oxides” from Live Demonstrations
42 “Disappearing Ink” from Live Demonstrations
43 “Metalloids” from Further Readings
Periodic Properties of the Elements 99
• Physical properties for the first transition-series elements are tabulated in Chapter 23 (section
23.1)
• Chapter 24 is devoted to the chemistry of carbon.
7.7 Trends for Group 1A and Group 2A Metals
• The alkali metals (group 1A) and the alkaline earth metals (group 2A) are often called the active
metals.
Group 1A: The Alkali Metals44,45,46,47,48,49,50,51,52,53
• The alkali metals are in Group 1A.
• Alkali metals are all soft.
• Their chemistry is dominated by the loss of their single s electron:
M  M+
+ e–
• Reactivity increases as we move down the group.
• Alkali metals react with hydrogen to form hydrides.
• In hydrides, the hydrogen is present as H–
, called the hydride ion.
2M(s) + H2(g)  2MH(s)
• Alkali metals react with water to form MOH and hydrogen gas:
2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
• Alkali metals produce different oxides when reacting with O2:
• 4Li(s) + O2(g)  2Li2O(s) (oxide)
• 2Na(s) + O2 (g)  Na2O2(s) (peroxide)
• K(s) + O2 (g)  KO2(s) (superoxide)
• Alkali metals emit characteristic colors when placed in a high-temperature flame.
• The s electron is excited by the flame and emits energy when it returns to the ground state.
• The Na line occurs at 589 nm (yellow), characteristic of the 3p  3s transition.
• The Li line is crimson red.
• The K line is lilac.
Group 2A: The Alkaline Earth Metals54,55,56,57

• Alkaline earth metals are harder and more dense than the alkali metals.
• Their chemistry is dominated by the loss of two s electrons:
M  M2+ + 2e–
Mg(s) + Cl2(g)  MgCl2(s)
2Mg(s) + O2(g)  2MgO(s)
44 Table 7.4
45 “Sodium and Potassium in Water” Movie from Instructor’s Resource CD/DVD
46 “Flame Tests for Metals” Movie from Instructor’s Resource CD/DVD
47 “A Dramatic Flame Test Demonstration” from Live Demonstrations
48 “Flame Tests for Metals” VCL Simulation from Instructor’s Resource CD/DVD
49 “Simple Flame Test Techniques Using Cotton Swabs” from Live Demonstrations
50 “Emission Spectra of Sodium and Mercury” VCL Simulation from Instructor’s Resource CD/DVD
51 “The Legend of Dr. Pepper/ Seven-Up” from Further Readings
52 “A Little Lithium May Be Just What the Doctor Ordered” from Further Readings
53 “A Variation on the Determination of the Properties of the Alkali Metals” from Further Readings
54 Table 7.5
55 “Producing Hydrogen Gas from Calcium Metal” from Live Demonstrations
56 “Life, Death, and Calcium” from Further Readings
57 “Update on Intake: Calcium Consumption Low” from Further Readings