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CHAPTER
8
Periodic Relationships
Among the Elements
INTRODUCTION
MANY
OF THE CHEMICAL PROPERTIES OF THE ELEMENTS CAN BE UN-
DERSTOOD IN TERMS OF THEIR ELECTRON CONFIGURATIONS.
8.1 DEVELOPMENT OF THE PERIODIC TABLE
BECAUSE
8.2 PERIODIC CLASSIFICATION OF THE
ELEMENTS
ELECTRONS FILL ATOMIC ORBITALS IN A FAIRLY REGULAR FASHION, IT IS
NOT SURPRISING THAT ELEMENTS WITH SIMILAR ELECTRON CONFIGURA-
8.3 PERIODIC VARIATION IN PHYSICAL
PROPERTIES
TIONS, SUCH AS SODIUM AND POTASSIUM, BEHAVE SIMILARLY IN MANY
8.4 IONIZATION ENERGY
RESPECTS AND THAT, IN GENERAL, THE PROPERTIES OF THE ELEMENTS
8.5 ELECTRON AFFINITY
EXHIBIT OBSERVABLE TRENDS.
8.6 VARIATION IN CHEMICAL PROPERTIES OF
THE REPRESENTATIVE ELEMENTS
CHEMISTS
IN THE NINETEENTH CENTURY
RECOGNIZED PERIODIC TRENDS IN THE PHYSICAL AND CHEMICAL PROPERTIES OF ELEMENTS LONG BEFORE QUANTUM THEORY CAME ONTO THE
SCENE.
ALTHOUGH
THESE CHEMISTS WERE NOT AWARE OF THE EXISTENCE
OF ELECTRONS AND PROTONS, THEIR EFFORTS TO SYSTEMATIZE THE CHEMISTRY OF THE ELEMENTS WERE REMARKABLY SUCCESSFUL.
THEIR
MAIN
SOURCES OF INFORMATION WERE THE ATOMIC MASSES OF THE ELEMENTS
AND OTHER KNOWN PHYSICAL AND CHEMICAL PROPERTIES.
287
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
8.1
DEVELOPMENT OF THE PERIODIC TABLE
In the nineteenth century, when chemists had only a vague idea of atoms and molecules and did not know of the existence of electrons and protons, they devised the periodic table using their knowledge of atomic masses. Accurate measurements of the
atomic masses of many elements had already been made. Arranging elements according to their atomic masses in a periodic table seemed logical to those chemists, who
felt that chemical behavior should somehow be related to atomic mass.
In 1864 the English chemist John Newlands noticed that when the elements were
arranged in order of atomic mass, every eighth element had similar properties. Newlands
referred to this peculiar relationship as the law of octaves. However, this law turned
out to be inadequate for elements beyond calcium, and Newlandss work was not accepted by the scientific community.
In 1869 the Russian chemist Dmitri Mendeleev and the German chemist Lothar
Meyer independently proposed a much more extensive tabulation of the elements based
on the regular, periodic recurrence of properties. Table 8.1 shows an early version of
Mendeleevs periodic table. Mendeleevs classification system was a great improvement over Newlandss for two reasons. First, it grouped the elements together more accurately, according to their properties. Equally important, it made possible the prediction of the properties of several elements that had not yet been discovered. For example,
Mendeleev proposed the existence of an unknown element that he called eka-aluminum
and predicted a number of its properties. (Eka is a Sanskrit word meaning first; thus
eka-aluminum would be the first element under aluminum in the same group.) When
gallium was discovered four years later, its properties matched the predicted properties of eka-aluminum remarkably well:
EKA-ALUMINUM (Ea)
Atomic mass
Melting point
Density
Formula of oxide
Appendix 1 explains the names and
symbols of the elements.
GALLIUM (Ga)
68 amu
Low
5.9 g/cm3
Ea2O3
69.9 amu
30.15C
5.94 g/cm3
Ga2O3
Mendeleevs periodic table included 66 known elements. By 1900, some 30 more had
been added to the list, filling in some of the empty spaces. Figure 8.1 charts the discovery of the elements chronologically.
Although this periodic table was a celebrated success, the early versions had some
glaring inconsistencies. For example, the atomic mass of argon (39.95 amu) is greater
than that of potassium (39.10 amu). If elements were arranged solely according to increasing atomic mass, argon would appear in the position occupied by potassium in
our modern periodic table (see the inside front cover). But no chemist would place argon, an inert gas, in the same group as lithium and sodium, two very reactive metals.
This and other discrepancies suggested that some fundamental property other than
John Alexander Reina Newlands (18381898). English chemist. Newlandss work was a step in the right direction in the
classification of the elements. Unfortunately, because of its shortcomings, he was subjected to much criticism, and even
ridicule. At one meeting he was asked if he had ever examined the elements according to the order of their initial letters!
Nevertheless, in 1887 Newlands was honored by the Royal Society of London for his contribution.
Dmitri Ivanovich Mendeleev (18361907). Russian chemist. His work on the periodic classification of elements is regarded
by many as the most significant achievement in chemistry in the nineteenth century.
Julius Lothar Meyer (18301895). German chemist. In addition to his contribution to the periodic table, Meyer also discovered the chemical affinity of hemoglobin for oxygen.
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8.1
TABLE 8.1
REIHEN
DEVELOPMENT OF THE PERIODIC TABLE
289
The Periodic Table as Drawn by Mendeleev*
GRUPPE I
R2O
H
7
GRUPPE II
RO
GRUPPE III
R2O3
GRUPPE IV
RH4
RO2
GRUPPE V
RH3
R2O5
GRUPPE VI
RH2
RO3
GRUPPE VII GRUPPE VIII
RH
R2O7
RO4
Be
B
C
N
O
F
1
2
Li
1
3
4
K
5
6
(Cu 63)
Zn 65
68
72
As 75
Se 78
Rb 85
Sr 87
?Yt 88
Zr 90
Nb 94
Mo 96
7
8
9
10
(Ag 108) Cd 112
In 113
Sn 118
Sb 122
Te 125
Cs 133
Ba 137
?Di 138 ?Ce 140
()
?Er 178 ?La 180 Ta 182
W 184
11
12
(Au
9,4
11
12
14
Na 23
Mg 24 Al 27,3
Si 28
39
Ca 40
44
Ti 48
V
199)
Hg
200
Ti
204
Pb 207
Th 231
16
P 31
S 32
51
Cr 52
Bi
208
U
240
Cl
Mn
19
35,5
55
Fe
56, Co 59,
Ni 59, Cu 63.
Br 80
100
Ru 104, Rh 104,
Pd 106, Ag 108.
J 127
Os 195, Ir 197,
Pt 198, Au 199.
*Spaces are left for the unknown elements with atomic masses 44, 68, 72, and 100.
FIGURE 8.1 A chronological
chart of the discovery of the elements. To date, 112 elements
have been identified.
120
Synthetic elements
100
Mendeleevs first
periodic table
Number of elements
80
60
Elements known prior to 1650:
40
Ag As Au C Cu Fe
Hg Pb S Sb Sn
20
0
1650
1700
1750
1800
1850
Year discovered
1900
1950
2000
atomic mass must be the basis of periodicity. This property turned out to be associated
with atomic number, a concept unknown to Mendeleev and his contemporaries.
Using data from -scattering experiments (see Section 2.2), Rutherford estimated
the number of positive charges in the nucleus of a few elements, but the significance
of these numbers was overlooked for several more years. In 1913 a young English
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290
PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
physicist, Henry Moseley, discovered a correlation between what he called atomic
number and the frequency of X rays generated by bombarding an element with highenergy electrons. Moseley noticed that the frequencies of X rays emitted from the elements could be correlated by the equation
v
a(Z
b)
where v is the frequency of the emitted X rays and a and b are constants that are the
same for all the elements. Thus, from the square root of the measured frequency of the
X rays emitted, we can determine the atomic number of the element.
With a few exceptions, Moseley found that atomic number increases in the same
order as atomic mass. For example, calcium is the twentieth element in order of increasing atomic mass, and it has an atomic number of 20. The discrepancies that had
puzzled earlier scientists now made sense. The atomic number of argon is 18 and that
of potassium is 19, so potassium should follow argon in the periodic table.
A modern periodic table usually shows the atomic number along with the element
symbol. As you already know, the atomic number also indicates the number of electrons in the atoms of an element. Electron configurations of elements help to explain
the recurrence of physical and chemical properties. The importance and usefulness of
the periodic table lie in the fact that we can use our understanding of the general properties and trends within a group or a period to predict with considerable accuracy the
properties of any element, even though that element may be unfamiliar to us.
8.2
PERIODIC CLASSIFICATION OF THE ELEMENTS
Figure 8.2 shows the periodic table together with the outermost ground-state electron
configurations of the elements. (The electron configurations of the elements are also
given in Table 7.3.) Starting with hydrogen, we see that subshells are filled in the order shown in Figure 7.23. According to the type of subshell being filled, the elements
can be divided into categories the representative elements, the noble gases, the transition elements (or transition metals), the lanthanides, and the actinides. The representative elements (also called main group elements) are the elements in Groups 1A
through 7A, all of which have incompletely filled s or p subshells of the highest principal quantum number. With the exception of helium, the noble gases (the Group 8A
elements) all have a completely filled p subshell. (The electron configurations are 1s2
for helium and ns2np6 for the other noble gases, where n is the principal quantum number for the outermost shell.)
The transition metals are the elements in Groups 1B and 3B through 8B, which
have incompletely filled d subshells, or readily produce cations with incompletely filled
d subshells. (These metals are sometimes referred to as the d-block transition elements.)
The nonsequential numbering of the transition metals in the periodic table (that is, 3B
8B, followed by 1B 2B) acknowledges a correspondence between the outer electron
configurations of these elements and those of the representative elements. For example, scandium and gallium both have 3 outer electrons. However, because they are in
different types of atomic orbitals, they are placed in different groups (3A and 3B). The
Henry Gwyn-Jeffreys Moseley (18871915). English physicist. Moseley discovered the relationship between X-ray spectra and atomic number. A lieutenant in the Royal Engineers, he was killed in action at the age of 28 during the British campaign in Gallipoli, Turkey.
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8.2
PERIODIC CLASSIFICATION OF THE ELEMENTS
1
1A
291
18
8A
1
1
H
1s1
2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
1s2
2
3
Li
2s1
4
Be
2s2
5
B
2s22p1
6
C
2s22p2
7
N
2s22p3
8
O
2s22p4
9
F
2s22p5
10
Ne
2s22p6
3
11
Na
3s1
12
Mg
3s2
3
3B
4
4B
5
5B
6
6B
7
7B
8
9
8B
10
11
1B
12
2B
13
Al
3s23p1
14
Si
3s23p2
15
P
3s23p3
16
S
3s23p4
17
Cl
3s23p5
18
Ar
3s23p6
4
19
K
4s1
20
Ca
4s2
21
Sc
4s23d1
22
Ti
4s23d2
23
V
4s23d 3
24
Cr
4s13d5
25
Mn
4s23d5
26
Fe
4s23d6
27
Co
4s23d 7
28
Ni
4s23d 8
29
Cu
4s13d10
30
Zn
4s23d10
31
Ga
4s24p1
32
Ge
4s24p2
33
As
4s24p3
34
Se
4s24p4
35
Br
4s24p5
36
Kr
4s24p6
5
37
Rb
5s1
38
Sr
5s2
39
Y
5s24d1
40
Zr
5s24d2
41
Nb
5s14d 4
42
Mo
5s14d5
43
Tc
5s24d5
44
Ru
5s14d 7
45
Rh
5s14d 8
46
Pd
4d10
47
Ag
5s14d10
48
Cd
5s24d10
49
In
5s25p1
50
Sn
5s25p2
51
Sb
5s25p3
52
Te
5s25p4
53
I
5s25p5
54
Xe
5s25p6
6
55
Cs
6s1
56
Ba
6s2
57
La
6s25d1
72
Hf
6s25d2
73
Ta
6s25d 3
74
W
6s25d4
75
Re
6s25d5
76
Os
6s25d6
77
Ir
6s25d 7
78
Pt
6s15d9
79
Au
6s15d10
80
Hg
6s25d10
81
Tl
6s26p1
82
Pb
6s26p2
83
Bi
6s26p3
84
Po
6s26p4
85
At
6s26p5
86
Rn
6s26p6
7
87
Fr
7s1
88
Ra
7s2
89
Ac
7s26d1
104
Rf
7s26d2
105
Ha
7s26d 3
106
Sg
7s26d4
107
Ns
7s26d5
108
Hs
7s26d6
109
Mt
7s26d 7
110
111
112
7s26d8
7s26d9
7s26d 10
58
Ce
6s24f 15d1
59
Pr
6s24f 3
60
Nd
6s24f 4
61
Pm
6s24f 5
62
Sm
6s24f 6
63
Eu
6s24f 7
64
Gd
6s24f 75d1
65
Tb
6s24f 9
66
Dy
6s24f 10
67
Ho
6s24f 11
68
Er
6s24f 12
69
Tm
6s24f 13
70
Yb
6s24f 14
71
Lu
6s24f 145d1
90
Th
7s26d2
91
Pa
7s25f 26d1
92
U
7s25f 36d1
93
Np
7s25f 46d1
94
Pu
7s25f 6
95
Am
7s25f 7
96
Cm
7s25f 76d1
97
Bk
7s25f 9
98
Cf
7s25f 10
99
Es
7s25f 11
100
Fm
7s25f 12
101
Md
7s25f 13
102
No
7s25f 14
103
Lr
7s25f 146d1
FIGURE 8.2 The groundstate electron configurations of
the elements. For simplicity,
only the configurations of the
outer electrons are shown.
TABLE 8.2 Electron
Configurations of Group 1A
and Group 2A Elements
GROUP 1A
Li
Na
K
Rb
Cs
Fr
Back
GROUP 2A
1
[He]2s
[Ne]3s1
[Ar]4s1
[Kr]5s1
[Xe]6s1
[Rn]7s1
Forward
Be
Mg
Ca
Sr
Ba
Ra
[He]2s2
[Ne]3s2
[Ar]4s2
[Kr]5s2
[Xe]6s2
[Rn]7s2
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metals iron (Fe), cobalt (Co), and nickel (Ni) do not fit this classification and are all
placed in Group 8B. The Group 2B elements, Zn, Cd, and Hg, are neither representative elements nor transition metals. There is no special name for this group of metals.
It should be noted that the designation of A and B groups is not universal. In Europe
the practice is to use B for representative elements and A for transition metals, which
is just the opposite of the American convention. The International Union of Pure and
Applied Chemistry (IUPAC) has recommended numbering the columns sequentially
with Arabic numerals 1 through 18 (see Figure 8.2). The proposal has sparked much
controversy in the international chemistry community, and its merits and drawbacks
will be deliberated for some time to come. In this text we will adhere to the American
designation.
The lanthanides and actinides are sometimes called f-block transition elements because they have incompletely filled f subshells. Figure 8.3 distinguishes the groups of
elements discussed here.
A clear pattern emerges when we examine the electron configurations of the elements in a particular group. The electron configurations for Groups 1A and 2A are
shown in Table 8.2. All members of the Group 1A alkali metals have similar outer electron configurations; each has a noble gas core and an ns1 outer electron. Similarly, the
Group 2A alkaline earth metals have a noble gas core and an outer electron configuration of ns2. The outer electrons of an atom, which are the ones involved in chemical
bonding, are often called valence electrons. The similarity of the outer electron configurations (that is, they have the same number and type of valence electrons) is what
makes the elements in the same group resemble one another in chemical behavior. This
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
Representative
elements
1
H
3
4
Li
Be
11
12
Na
Mg
3
3B
Lanthanides
13
3A
Transition
metals
2
2A
Zinc
Cadium
Mercury
Noble gases
1
1A
Actinides
5
6
7
8
9
10
B
C
N
O
F
Ne
13
14
15
16
17
18
Al
Si
P
S
Cl
Ar
4
4B
5
5B
6
6B
7
7B
8
9
8B
18
8A
10
11
1B
12
2B
14
4A
15
5A
16
6A
17
7A
2
He
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
110
111
112
87
88
89
104
105
106
107
108
109
Fr
Ra
Ac
Rf
Ha
Sg
Ns
Hs
Mt
FIGURE 8.3 Classification
of the elements. Note that
the Group 2B elements are
often classified as transition
metals even though they do
not exhibit the characteristics of the transition metals.
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
observation holds true for the other representative elements. Thus, for instance, the
halogens (the Group 7A elements) all have outer electron configurations of ns2np5, and
they have very similar properties. We must be careful, however, in predicting properties for Groups 3A through 7A. For example, the elements in Group 4A all have the
same outer electron configuration, ns2np2, but there is some variation in chemical properties among these elements: Carbon is a nonmetal, silicon and germanium are metalloids, and tin and lead are metals.
As a group, the noble gases behave very similarly. With the exception of krypton
and xenon, the rest of these elements are totally inert chemically. The reason is that
these elements all have completely filled outer ns and np subshells, a condition that
represents great stability. Although the outer electron configuration of the transition
metals is not always the same within a group and there is no regular pattern in the
change of the electron configuration from one metal to the next in the same period, all
transition metals share many characteristics that set them apart from other elements.
The reason is that these metals all have an incompletely filled d subshell. Likewise,
the lanthanide (and the actinide) elements resemble one another because they have incompletely filled f subshells.
EXAMPLE 8.1
A neutral atom of a certain element has 15 electrons. Without consulting a periodic
table, answer the following questions: (a) What is the ground-state electron configuration of the element? (b) How should the element be classified? (c) Are the atoms
of this element diamagnetic or paramagnetic?
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PERIODIC CLASSIFICATION OF THE ELEMENTS
293
(a) Using the building-up principle and knowing the maximum capacity
of s and p subshells, we can write the ground-state electron configuration of the element as 1s22s22p63s23p3.
(b) Since the 3p subshell is not completely filled, this is a representative element.
Based on the information given, we cannot say whether it is a metal, a nonmetal,
or a metalloid.
(c) According to Hunds rule, the three electrons in the 3p orbitals have parallel
spins. Therefore, the atoms of this element are paramagnetic, with three unpaired
spins. (Remember, we saw in Chapter 7 that any atom that contains an odd number
of electrons must be paramagnetic.)
Answer
Similar problem: 8.20.
PRACTICE EXERCISE
A neutral atom of a certain element has 20 electrons. (a) Write the ground-state electron configuration of the element, (b) classify the element, and (c) determine whether
the atoms of this element are diamagnetic or paramagnetic.
REPRESENTING FREE ELEMENTS IN CHEMICAL EQUATIONS
Having classified the elements according to their ground-state electron configurations,
we can now look at the way chemists represent metals, metalloids, and nonmetals as
free elements in chemical equations. Because metals do not exist in discrete molecular units, we always use their empirical formulas in chemical equations. The empirical
formulas are the same as the symbols that represent the elements. For example, the
empirical formula for iron is Fe, the same as the symbol for the element.
For nonmetals there is no single rule. Carbon, for example, exists as an extensive
three-dimensional network of atoms, and so we use its empirical formula (C) to represent elemental carbon in chemical equations. But hydrogen, nitrogen, oxygen, and
the halogens exist as diatomic molecules, and so we use their molecular formulas (H2,
N2, O2, F2, Cl2, Br2, I2) in equations. The stable form of phosphorus is molecular (P4),
and so we use P4. For sulfur chemists often use the empirical formula (S) in chemical
equations, rather than S8 which is the stable form. Thus, instead of writing the equation for the combustion of sulfur as
S8(s)
Note that these two equations for
the combustion of sulfur have identical stoichiometry. This correspondence should not be surprising,
since both equations describe the
same chemical system. In both cases
a number of sulfur atoms react with
twice as many oxygen atoms.
8O2(g) 88n 8SO2(g)
we usually write
S(s)
O2(g) 88n SO2(g)
All the noble gases are monatomic species; thus we use their symbols: He, Ne, Ar, Kr,
Xe, and Rn. The metalloids, like the metals, all have complex three-dimensional networks, and we represent them, too, with their empirical formulas, that is, their symbols: B, Si, Ge, and so on.
ELECTRON CONFIGURATIONS OF CATIONS AND ANIONS
Because many ionic compounds are made up of monatomic anions and cations, it is
helpful to know how to write the electron configurations of these ionic species. Just as
for neutral atoms, we use the Pauli exclusion principle and Hunds rule in writing the
ground-state electron configurations of cations and anions. We will group the ions in
two categories for discussion.
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
Ions Derived from Representative Elements
Ions formed from neutral atoms of most representative elements have the noble-gas
outer-electron configuration of ns2np6. In the formation of a cation from the neutral
atom of a representative element, one or more electrons are removed from the highest
occupied n shell. Following are the electron configurations of some neutral atoms and
their corresponding cations:
Na: [Ne]3s1
Ca: [Ar]4s2
Al: [Ne]3s23p1
Na : [Ne]
Ca2 : [Ar]
Al3 : [Ne]
Note that each ion has a stable noble gas configuration.
In the formation of an anion, one or more electrons are added to the highest partially filled n shell. Consider the following examples:
H:
F:
O:
N:
1s1
1s22s22p5
1s22s22p4
1s22s22p3
H
F
O2
N3
:
:
:
:
1s2 or [He]
1s22s22p6 or [Ne]
1s22s22p6 or [Ne]
1s22s22p6 or [Ne]
All of these anions also have stable noble gas configurations. Notice that F , Na , and
Ne (and Al3 , O2 , and N3 ) have the same electron configuration. They are said to
be isoelectronic because they have the same number of electrons, and hence the same
ground-state electron configuration. Thus H and He are also isoelectronic.
Cations Derived from Transition Metals
Bear in mind that the order of electron filling does not determine or
predict the order of electron removal
for transition metals.
8.3
In Section 7.10 we saw that in the first-row transition metals (Sc to Cu), the 4s orbital
is always filled before the 3d orbitals. Consider manganese, whose electron configuration is [Ar]4s23d 5. When the Mn2 ion is formed, we might expect the two electrons
to be removed from the 3d orbitals to yield [Ar]4s23d 3. In fact, the electron configuration of Mn2 is [Ar]3d 5! The reason is that the electron-electron and electronnucleus interactions in a neutral atom can be quite different from those in its ion. Thus,
whereas the 4s orbital is always filled before the 3d orbital in Mn, electrons are removed from the 4s orbital in forming Mn2 because the 3d orbital is more stable than
the 4s orbital in transition metal ions. Therefore, when a cation is formed from an atom
of a transition metal, electrons are always removed first from the ns orbital and then
from the (n 1)d orbitals.
Keep in mind that most transition metals can form more than one cation and that
frequently the cations are not isoelectronic with the preceding noble gases.
PERIODIC VARIATION IN PHYSICAL PROPERTIES
As we have seen, the electron configurations of the elements show a periodic variation
with increasing atomic number. Consequently, there are also periodic variations in physical and chemical behavior. In this section and the next two, we will examine some
physical properties of elements that are in the same group or period and additional
properties that influence the chemical behavior of the elements. First, lets look at the
concept of effective nuclear charge, which has a direct bearing on atomic size and on
the tendency for ionization.
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295
EFFECTIVE NUCLEAR CHARGE
In Chapter 7, we discussed the shielding effect that electrons close to the nucleus have
on outer-shell electrons in many-electron atoms. The presence of shielding electrons
reduces the electrostatic attraction between the positively charged protons in the nucleus and the outer electrons. Moreover, the repulsive forces between electrons in a
many-electron atom further offset the attractive force exerted by the nucleus. The concept of effective nuclear charge allows us to account for the effects of shielding on periodic properties.
Consider, for example, the helium atom, which has the ground-state electron configuration 1s2. Heliums two protons give the nucleus a charge of 2, but the full attractive force of this charge on the two 1s electrons is partially offset by electron-electron repulsion. Consequently we say that the 1s electrons shield each other from the
nucleus. The effective nuclear charge (Zeff), which is the charge felt by an electron, is
given by
Zeff
Z
where Z is the actual nuclear charge (that is, the atomic number of the element) and
(sigma) is called the shielding constant (also called the screening constant). The shielding constant is greater than zero but smaller than Z.
One way to illustrate electron shielding is to consider the amounts of energy required to remove the two electrons from a helium atom. Measurements show that it
takes 2373 kJ of energy to remove the first electron from 1 mole of He atoms and
5251 kJ of energy to remove the remaining electron from 1 mole of He ions. The reason it takes so much more energy to remove the second electron is that with only one
electron present, there is no shielding, and the electron feels the full effect of the 2
nuclear charge.
For atoms with three or more electrons, the electrons in a given shell are shielded
by electrons in inner shells (that is, shells closer to the nucleus) but not by electrons
in outer shells. Thus, in a neutral lithium atom, whose electron configuration is 1s22s1,
the 2s electron is shielded by the two 1s electrons, but the 2s electron does not have a
shielding effect on the 1s electrons. In addition, filled inner shells shield outer electrons more effectively than electrons in the same subshell shield each other.
(a)
ATOMIC RADIUS
(b)
FIGURE 8.4 (a) In metals such
as beryllium, the atomic radius is
defined as one-half the distance
between the centers of two adjacent atoms. (b) For elements that
exist as diatomic molecules, such
as iodine, the radius of the atom
is defined as one-half the distance between the centers of the
atoms in the molecule.
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A number of physical properties, including density, melting point, and boiling point,
are related to the sizes of atoms, but atomic size is difficult to define. As we saw in
Chapter 7, the electron density in an atom extends far beyond the nucleus, but we normally think of atomic size as the volume containing about 90 percent of the total electron density around the nucleus. When we must be even more specific, we define the
size of an atom in terms of its atomic radius, which is one-half the distance between
the two nuclei in two adjacent metal atoms.
For atoms linked together to form an extensive three-dimensional network, atomic
radius is simply one-half the distance between the nuclei in two neighboring atoms
[Figure 8.4(a)]. For elements that exist as simple diatomic molecules, the atomic radius is one-half the distance between the nuclei of the two atoms in a particular molecule [Figure 8.4(b)].
Figure 8.5 shows the atomic radii of many elements according to their positions
in the periodic table, and Figure 8.6 plots the atomic radii of these elements against
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.5 Atomic radii (in
picometers) of representative elements according to their positions
in the periodic table. Note that
there is no general agreement on
the size of atomic radii. We focus only on the trends in atomic
radii, not on their precise values.
Increasing atomic radius
1A
2A
3A
4A
5A
6A
7A
8A
H
He
32
50
B
C
N
O
F
Ne
112
98
91
92
73
72
70
Na
Mg
Al
Si
P
S
Cl
Ar
186
160
143
132
128
127
99
98
K
Ca
Ga
Ge
As
Se
Br
Kr
227
197
135
137
139
140
114
112
Rb
Sr
In
Sn
Sb
Te
I
Xe
248
215
166
162
159
160
133
131
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
265
222
171
175
170
164
142
140
Be
152
Increasing atomic radius
Li
their atomic numbers. Periodic trends are clearly evident. In studying the trends, bear
in mind that the atomic radius is determined to a large extent by the strength of the attraction between the outer-shell electrons and the nucleus. The larger the effective nuclear charge, the stronger the hold of the nucleus on these electrons, and the smaller
the atomic radius. Consider the second-period elements from Li to F, for example.
Moving from left to right, we find that the number of electrons in the inner shell (1s2)
remains constant while the nuclear charge increases. The electrons that are added to
counterbalance the increasing nuclear charge are ineffective in shielding one another.
Consequently, the effective nuclear charge increases steadily while the principal quantum number remains constant (n 2). For example, the outer 2s electron in lithium is
shielded from the nucleus (which has 3 protons) by the two 1s electrons. As an approximation, we assume that the shielding effect of the two 1s electrons is to cancel
two positive charges in the nucleus. Thus the 2s electron only feels the attraction of
one proton in the nucleus; the effective nuclear charge is 1. In beryllium (1s22s2),
each of the 2s electrons is shielded by the inner two 1s electrons, which cancel two of
the four positive charges in the nucleus. Because the 2s electrons do not shield each
other as effectively, the net result is that the effective nuclear charge of each 2s electron is greater than 1. Thus as the effective nuclear charge increases, the atomic radius decreases steadily from lithium to fluorine.
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8.3
FIGURE 8.6 Plot of atomic
radii (in picometers) of elements
against their atomic numbers.
297
PERIODIC VARIATION IN PHYSICAL PROPERTIES
300
Cs
Rb
250
K
Atomic radius (pm)
200
Na
Li
Po
150
I
Br
100
Cl
F
50
0
10
20
30
40
50
Atomic number
60
70
80
90
Within a group of elements we find that atomic radius increases with increasing
atomic number. For the alkali metals in Group 1A, the outermost electron resides in
the ns orbital. Since orbital size increases with the increasing principal quantum number n, the size of the metal atoms increases from Li to Cs. We can apply the same reasoning to the elements in other groups.
EXAMPLE 8.2
Referring to a periodic table, arrange the following atoms in order of increasing radius: P, Si, N.
Note that N and P are in the same group (Group 5A) and that N is above
P. Therefore, the radius of N is smaller than that of P (atomic radius increases as
we go down a group). Both Si and P are in the third period, and Si is to the left of
P. Therefore, the radius of P is smaller than that of Si (atomic radius decreases as
we move from left to right across a period). Thus the order of increasing radius is
N P Si.
Answer
Similar problems: 8.37, 8.38.
PRACTICE EXERCISE
Arrange the following atoms in order of decreasing radius: C, Li, Be.
IONIC RADIUS
Ionic radius is the radius of a cation or an anion. It can be measured by X-ray diffraction (see Chapter 11). Ionic radius affects the physical and chemical properties of
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.7 Comparison of
atomic radii with ionic radii.
(a) Alkali metals and alkali metal
cations. (b) Halogens and halide
ions.
300
300
Cs
Rb
250
250
K
I
200
Na
Cl
Li
Cs+
150
K+
Rb+
Radius (pm)
Radius (pm)
200
Br
100
150
F
I
Br
100
Cl
Na+
50
0
10
F
Li+
50
20 30 40 50
Atomic number
(a)
60
0
10
20 30 40 50
Atomic number
(b)
60
an ionic compound. For example, the three-dimensional structure of an ionic compound
depends on the relative sizes of its cations and anions.
When a neutral atom is converted to an ion, we expect a change in size. If the
atom forms an anion, its size (or radius) increases, since the nuclear charge remains
the same but the repulsion resulting from the additional electron(s) enlarges the domain of the electron cloud. On the other hand, removing one or more electrons from
an atom reduces electron-electron repulsion but the nuclear charge remains the same,
so the electron cloud shrinks, and the cation is smaller than the atom. Figure 8.7 shows
the changes in size that result when alkali metals are converted to cations and halogens are converted to anions; Figure 8.8 shows the changes in size that occur when a
lithium atom reacts with a fluorine atom to form a LiF unit.
Figure 8.9 shows the radii of ions derived from the familiar elements, arranged
according to elements positions in the periodic table. We can see parallel trends between atomic radii and ionic radii. For example, from top to bottom both the atomic
radius and the ionic radius increase within a group. For ions derived from elements in
different groups, a size comparison is meaningful only if the ions are isoelectronic. If
we examine isoelectronic ions, we find that cations are smaller than anions. For example, Na is smaller than F . Both ions have the same number of electrons, but Na
FIGURE 8.8 Changes in the
sizes of Li and F when they react
to form LiF.
+
Li
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Li+
PERIODIC VARIATION IN PHYSICAL PROPERTIES
299
Be2+
N3
65
K+
Ca2+
133
99
140
136
S2
Cl
184
Mg2+
95
171
181
Se2
Br
195
Te2
I
31
Na+
F
198
60
O2
Al3+
Fe3+
Ti4+
V5+
Sc3+
Rb+
148
Fe2+
Cr3+
50
Ni2+
Mn2+
Co2+
Zn2+
Ga3+
Cu+
81
68
59 64
80
Sr2+
60
77
72
69
96
62
Sb5+
In3+ Sn4+
Ag+
126
113
74
Cd2+
97
81
71
Pb4+
Cs+
Ba2+
169
135
Au+
Hg2+
Tl3+
62
221
216
137
110
95
84
FIGURE 8.9 The radii (in picometers) of ions of familiar elements arranged according to the elements positions in the periodic table.
(Z 11) has more protons than F (Z 9). The larger effective nuclear charge of Na
results in a smaller radius.
Focusing on isoelectronic cations, we see that the radii of tripositive ions (ions
that bear three positive charges) are smaller than those of dipositive ions (ions that bear
two positive charges), which in turn are smaller than unipositive ions (ions that bear
one positive charge). This trend is nicely illustrated by the sizes of three isoelectronic
ions in the third period: Al3 , Mg2 , and Na (see Figure 8.9). The Al3 ion has the
same number of electrons as Mg2 , but it has one more proton. Thus the electron cloud
in Al3 is pulled inward more than that in Mg2 . The smaller radius of Mg2 compared with that of Na can be similarly explained. Turning to isoelectronic anions, we
find that the radius increases as we go from ions with uninegative charge ( ) to those
with dinegative charge (2 ), and so on. Thus the oxide ion is larger than the fluoride
ion because oxygen has one fewer proton than fluorine; the electron cloud is spread
out more in O2 .
EXAMPLE 8.3
For each of the following pairs, indicate which one of the two species is larger:
(a) N3 or F ; (b) Mg2 or Ca2 ; (c) Fe2 or Fe3 .
(a) N3 and F are isoelectronic anions. Since N3 has only seven protons and F has nine, N3 is larger.
(b) Both Mg and Ca belong to Group 2A (the alkaline metals). The Ca2 ion is
larger than Mg2 because Cas valence electrons are in a larger shell (n 4) than
are Mgs (n 3).
Answer
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.10 The third-period
elements. The photograph of argon, which is a colorless, odorless gas, shows the color emitted
by the gas from a discharge
tube.
Magnesium (Mg)
Sodium (Na)
Silicon (Si)
Aluminum (Al)
Phosphorous (P)
Chlorine (Cl2)
Sulfur (S)
Similar problems: 8.43, 8.45.
Argon (Ar)
(c) Both ions have the same nuclear charge, but Fe2 has one more electron and
hence greater electron-electron repulsion. The radius of Fe2 is larger.
PRACTICE EXERCISE
Select the smaller ion in each of the following pairs: (a) K , Li ; (b) Au , Au3 ;
(c) P3 , N3 .
VARIATION OF PHYSICAL PROPERTIES ACROSS A PERIOD AND WITHIN A GROUP
From left to right across a period there is a transition from metals to metalloids to nonmetals. Consider the third-period elements from sodium to argon (Figure 8.10). Sodium,
the first element in the third period, is a very reactive metal, whereas chlorine, the sec-
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IONIZATION ENERGY
301
ond-to-last element of that period, is a very reactive nonmetal. In between, the elements show a gradual transition from metallic properties to nonmetallic properties.
Sodium, magnesium, and aluminum all have extensive three-dimensional atomic networks, which are held together by forces characteristic of the metallic state. Silicon is
a metalloid; it has a giant three-dimensional structure in which the Si atoms are held
together very strongly. Starting with phosphorus, the elements exist in simple, discrete
molecular units (P4, S8, Cl2, and Ar) which have low melting points and boiling points.
Within a periodic group the physical properties vary more predictably, especially
if the elements are in the same physical state. For example, the melting points of argon and xenon are 189.2C and 111.9C, respectively. We can estimate the melting point of the intermediate element krypton by taking the average of these two values as follows:
melting point of Kr
[( 189.2C)
( 111.9C)]
2
150.6C
This value is quite close to the actual melting point of 156.6C.
The Chemistry in Action essay on p. 302 illustrates one interesting application of
periodic group properties.
8.4
IONIZATION ENERGY
Not only is there a correlation between electron configuration and physical properties,
but a close correlation also exists between electron configuration (a microscopic property) and chemical behavior (a macroscopic property). As we will see throughout this
book, the chemical properties of any atom are determined by the configuration of the
atoms valence electrons. The stability of these outermost electrons is reflected directly
in the atoms ionization energies. Ionization energy is the minimum energy (in kJ/mol)
required to remove an electron from a gaseous atom in its ground state. In other words,
ionization energy is the amount of energy in kilojoules needed to strip one mole of
electrons from one mole of gaseous atoms. Gaseous atoms are specified in this definition because an atom in the gas phase is virtually uninfluenced by its neighbors and
so there are no intermolecular forces (that is, forces between molecules) to take into
account when measuring ionization energy.
The magnitude of ionization energy is a measure of how tightly the electron is
held in the atom. The higher the ionization energy, the more difficult it is to remove
the electron. For a many-electron atom, the amount of energy required to remove the
first electron from the atom in its ground state,
energy
X(g) 88n X (g)
e
(8.1)
is called the first ionization energy (I1). In equation (8.1), X represents an atom of any
element, e is an electron, and g denotes the gaseous state. The second ionization energy (I2) and the third ionization energy (I3) are shown in the following equations:
energy
energy
X (g) 88n X2 (g)
2
3
X (g) 88n X (g)
e
second ionization
e
third ionization
The pattern continues for the removal of subsequent electrons.
When an electron is removed from a neutral atom, the repulsion among the remaining electrons decreases. Since the nuclear charge remains constant, more energy
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
Back
The Third Liquid Element?
Of the 112 known elements, 11 are gases under atmospheric conditions. Six of these are the Group 8A
elements (the noble gases He, Ne, Ar, Kr, Xe, and
Rn), and the other five are hydrogen (H2), nitrogen
(N2), oxygen (O2), fluorine (F2), and chlorine (Cl2).
Curiously, only two elements are liquids at 25C: mercury (Hg) and bromine (Br2).
We do not know the properties of all the known
elements because some of them have never been prepared in quantities large enough for investigation. In
these cases we must rely on periodic trends to predict
their properties. What are the chances, then, of discovering a third liquid element?
Let us look at francium (Fr), the last member of
Group 1A, to see if it might be a liquid at 25C. All
of franciums isotopes are radioactive. The most stable isotope is francium-223, which has a half-life of
21 minutes. (Half-life is the time it takes for one-half
of the nuclei in any given amount of a radioactive sub-
stance to disintegrate.) This short half-life means that
only very small traces of francium could possibly exist on Earth. And although it is feasible to prepare
francium in the laboratory, no weighable quantity of
the element has been prepared or isolated. Thus we
know very little about franciums physical and chemical properties. Yet we can use the group periodic
trends to predict some of those properties.
Take franciums melting point as an example. The
plot shows how the melting points of the alkali metals
vary with atomic number. From lithium to sodium, the
melting point drops 81.4; from sodium to potassium,
34.6; from potassium to rubidium, 24; from rubidium to cesium, 11. On the basis of this trend, we can
predict that the change from cesium to francium would
be about 5. If so, the melting point of francium would
be 23C, which would make it a liquid under atmospheric conditions.
180
Li
150
Melting point (C)
Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action
Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry
120
90
Na
K
60
Rb
Cs
30
40
60
80
100
Atomic number
A plot of the melting points of the alkali metals versus
their atomic numbers. By extrapolation, the melting point
of francium should be 23C.
0
False-colored image of francium-210 atoms created by
bombarding gold with oxygen and trapped by laser
beams. The central spot is 1 mm in diameter and consists of about 10,000 atoms. The yellow light is the fluorescence of the Fr atoms induced by laser. The other
bright areas come from laser light scattered off the glass
surfaces.
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8.4
TABLE 8.3
IONIZATION ENERGY
303
The Ionization Energies (kJ/mol) of the First 20 Elements
Z
ELEMENT
FIRST
SECOND
THIRD
FOURTH
FIFTH
SIXTH
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
1312
2373
520
899
801
1086
1400
1314
1680
2080
495.9
738.1
577.9
786.3
1012
999.5
1251
1521
418.7
589.5
5251
7300
1757
2430
2350
2860
3390
3370
3950
4560
1450
1820
1580
1904
2250
2297
2666
3052
1145
11815
14850
3660
4620
4580
5300
6050
6120
6900
7730
2750
3230
2910
3360
3820
3900
4410
4900
21005
25000
6220
7500
7470
8400
9370
9540
10500
11600
4360
4960
4660
5160
5770
5900
6500
32820
38000
9400
11000
11000
12200
13400
13600
14800
16000
6240
6990
6540
7240
8000
8100
47261
53000
13000
15200
15000
16600
18000
18400
20000
21000
8500
9300
8800
9600
11000
is needed to remove another electron from the positively charged ion. Thus, ionization
energies always increase in the following order:
I1
I2
I3
...
Table 8.3 lists the ionization energies of the first 20 elements. Ionization is always an
endothermic process. By convention, energy absorbed by atoms (or ions) in the ionization process has a positive value. Thus ionization energies are all positive quantities. Figure 8.11 shows the variation of the first ionization energy with atomic number. The plot clearly exhibits the periodicity in the stability of the most loosely held
electron. Note that, apart from small irregularities, the first ionization energies of elements in a period increase with increasing atomic number. This trend is due to the increase in effective nuclear charge from left to right (as in the case of atomic radii variation). A larger effective nuclear charge means a more tightly held outer electron, and
hence a higher first ionization energy. A notable feature of Figure 8.11 is the peaks,
which correspond to the noble gases. The high ionization energies of the noble gases,
stemming from their stable ground-state electron configurations, account for the fact
that most of them are chemically unreactive. In fact, helium (1s2) has the highest first
ionization energy of all the elements.
At the bottom of the graph in Figure 8.11 are the Group 1A elements (the alkali
metals) which have the lowest first ionization energies. Each of these metals has one
valence electron (the outermost electron configuration is ns1) which is effectively
shielded by the completely filled inner shells. Consequently, it is energetically easy to
remove an electron from the atom of an alkali metal to form a unipositive ion (Li ,
Na , K , . . .). Significantly, the electron configurations of these cations are isoelectronic with those noble gases just preceding them in the periodic table.
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.11 Variation of the
first ionization energy with
atomic number. Note that the noble gases have high ionization
energies, whereas the alkali metals and alkaline earth metals
have low ionization energies.
2500
He
Ne
First ionization energy (kJ/mol)
304
2000
Ar
Kr
1500
Xe
H
Rn
1000
500
Li
0
Na
10
Rb
K
20
30
Cs
40
50
Atomic number (Z )
60
70
80
90
The group 2A elements (the alkaline earth metals) have higher first ionization energies than the alkali metals do. The alkaline earth metals have two valence electrons
(the outermost electron configuration is ns2). Because these two s electrons do not
shield each other well, the effective nuclear charge for an alkaline earth metal atom is
larger than that for the preceding alkali metal. Most alkaline earth compounds contain
dipositive ions (Mg2 , Ca2 , Sr2 , Ba2 ). The Be2 ion is isoelectronic with Li and
with He, Mg2 is isoelectronic with Na and with Ne, and so on.
As Figure 8.11 shows, metals have relatively low ionization energies compared to
nonmetals. The ionization energies of the metalloids generally fall between those of
metals and nonmetals. The difference in ionization energies suggests why metals always form cations and nonmetals form anions in ionic compounds. (The only important nonmetallic cation is the ammonium ion, NH4 .) For a given group, ionization energy decreases with increasing atomic number (that is, as we move down the group).
Elements in the same group have similar outer electron configurations. However, as
the principal quantum number n increases, so does the average distance of a valence
electron from the nucleus. A greater separation between the electron and the nucleus
means a weaker attraction, so that it becomes increasingly easier to remove the first
electron as we go from element to element down a group. Thus the metallic character
of the elements within a group increases from top to bottom. This trend is particularly
noticeable for elements in Groups 3A to 7A. For example, in Group 4A, carbon is a
nonmetal, silicon and germanium are metalloids, and tin and lead are metals.
Although the general trend in the periodic table is for first ionization energies to
increase from left to right, some irregularities do exist. The first exception occurs between Group 2A and 3A elements in the same period (for example, between Be and
B and between Mg and Al).The Group 3A elements have lower first ionization energies than 2A elements because they all have a single electron in the outermost p subshell (ns2np1), which is well shielded by the inner electrons and the ns2 electrons.
Therefore, less energy is needed to remove a single p electron than to remove a paired
s electron from the same principal energy level. The second irregularity occurs between
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ELECTRON AFFINITY
305
Groups 5A and 6A (for example, between N and O and between P and S). In the Group
5A elements (ns2np3) the p electrons are in three separate orbitals according to Hunds
rule. In Group 6A (ns2np4) the additional electron must be paired with one of the three
p electrons. The proximity of two electrons in the same orbital results in greater electrostatic repulsion, which makes it easier to ionize an atom of the Group 6A element,
even though the nuclear charge has increased by one unit. Thus the ionization energies
for Group 6A elements are lower than those for Group 5A elements in the same period.
The following example compares the ionization energies of some elements.
EXAMPLE 8.4
(a) Which atom should have a smaller first ionization energy: oxygen or sulfur?
(b) Which atom should have a higher second ionization energy: lithium or beryllium?
(a) Oxygen and sulfur are members of Group 6A. They have the same valence electron configuration (ns2np4), but the 3p electron in sulfur is farther from
the nucleus and experiences less nuclear attraction than the 2p electron in oxygen.
Thus, following the general rule that the ionization energy of elements decreases as
we move down a periodic group, we predict that sulfur should have a smaller first
ionization energy. Table 8.3 confirms our reasoning.
(b) The electron configurations of Li and Be are 1s22s1 and 1s22s2, respectively. For
the second ionization process we write
Answer
Li (g) 88n Li2 (g)
1s2
1s1
Be (g) 88n Be2 (g)
1s22s1
1s2
Similar problem: 8.53.
Since 1s electrons shield 2s electrons much more effectively than they shield each
other, we predict that it should be much easier to remove a 2s electron from Be
than to remove a 1s electron from Li . This is consistent with the data in Table 8.3.
PRACTICE EXERCISE
(a) Which of the following atoms should have a larger first ionization energy: N or
P? (b) Which of the following atoms should have a smaller second ionization energy: Na or Mg?
8.5
ELECTRON AFFINITY
Another property that greatly influences the chemical behavior of atoms is their ability to accept one or more electrons. This property is called electron affinity, which is
the energy change that occurs when an electron is accepted by an atom in the gaseous
state to form an anion. The equation is
X(g)
e 88n X (g)
(8.2)
The sign of the electron affinity is opposite to the one we use for ionization energy. As
we saw in Section 8.5, a positive ionization energy means that energy must be supplied
to remove an electron. A positive electron affinity, on the other hand, signifies that en-
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
TABLE 8.4 Electron Affinities (kJ/mol) of Some Representative
Elements and the Noble Gases*
1A
H
73
Li
60
Na
53
K
48
Rb
47
Cs
45
2A
Be
0
Mg
0
Ca
2.4
Sr
4.7
Ba
14
3A
B
27
Al
44
Ga
29
In
29
Tl
30
4A
5A
C
122
Si
134
Ge
118
Sn
121
Pb
110
N
0
P
72
As
77
Sb
101
Bi
110
6A
7A
F
328
Cl
349
Br
325
I
295
At
?
O
141
S
200
Se
195
Te
190
Po
?
8A
He
0
Ne
0
Ar
0
Kr
0
Xe
0
Rn
0
*The electron affinities of the noble gases, Be, and Mg have not been determined experimentally,
but are believed to be close to zero or negative.
ergy is liberated when an electron is added to an atom. To clarify this apparent paradox, let us consider the process in which a gaseous fluorine atom accepts an electron:
F(g)
e 88n F (g)
H
328 kJ/mol
The sign of the enthalpy change indicates that this is an exothermic process; however,
the electron affinity of fluorine is assigned a value of 328 kJ/mol. Thus we can think
of electron affinity as the energy that must be supplied to remove an electron from a
negative ion. For the removal of an electron from a fluoride ion, we have
F (g) 88n F(g)
e
H
328 kJ/mol
Two features of electron affinity to remember are: (1) The electron affinity of an element is equal to the enthalpy change that accompanies the ionization process of its anion, and (2) a large positive electron affinity means that the negative ion is very stable
(that is, the atom has a great tendency to accept an electron), just as a high ionization
energy of an atom means that the atom is very stable.
Experimentally, electron affinity is determined by removing the additional electron
from an anion. In contrast to ionization energies, however, electron affinities are difficult to measure because the anions of many elements are unstable. Table 8.4 shows the
electron affinities of some representative elements and the noble gases, and Figure 8.12
plots the electron affinities of the first 20 elements versus atomic number. The overall
trend is an increase in the tendency to accept electrons (electron affinity values become
more positive) from left to right across a period. The electron affinities of metals are
generally lower than those of nonmetals. The values vary little within a given group.
The halogens (Group 7A) have the highest electron affinity values. This is not surprising when we realize that by accepting an electron, each halogen atom assumes the stable electron configuration of the noble gas immediately to its right. For example, the
electron configuration of F is 1s22s22p6, or [Ne]; for Cl it is [Ne]3s23p6 or [Ar]; and
so on. Calculations show that the noble gases all have electron affinities of less than
zero. Thus the anions of these gases, if formed, would be inherently unstable.
The electron affinity of oxygen has a positive value (141 kJ/mol), which means
that the process
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FIGURE 8.12 A plot of electron affinity against atomic number for the first 20 elements.
307
ELECTRON AFFINITY
400
Cl
F
Br
I
Electron affinity (kJ/mol)
300
200
Si
Sn
Ge
C
100
Li
Na
0
10
K
20
Rb
30
Atomic number (Z )
O(g)
Cs
40
50
e 88n O (g)
60
H
141 kJ
is favorable (exothermic). On the other hand, the electron affinity of the O
highly negative ( 780 kJ/mol), which means the process
O (g)
e 88n O2 (g)
H
ion is
780 kJ
is endothermic even though the O2 ion is isoelectronic with the noble gas Ne. This
process is unfavorable in the gas phase because the resulting increase in electron-electron repulsion outweighs the stability gained by achieving a noble gas configuration.
However, note that O2 is common in ionic compounds (for example, Li2O and MgO);
in solids, the O2 ion is stabilized by the neighboring cations. We will study the stability of ionic compounds in the next chapter.
The following example shows why the alkaline earth metals do not have a great
tendency to accept electrons.
EXAMPLE 8.5
Why are the electron affinities of the alkaline earth metals, shown in Table 8.4, either negative or small positive values?
Answer
The valence configuration of the alkaline earth metals is ns2. For the
process
M(g)
e 88n M (g)
where M denotes a member of the Group 2A family, the extra electron must enter
the np subshell, which is effectively shielded by the two ns electrons (the np elec-
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
Similar problem: 8.61.
trons are farther away from the nucleus than the ns electrons) and the inner electrons. Consequently, alkaline earth metals have little tendency to pick up an extra
electron.
PRACTICE EXERCISE
Is it likely that Ar will form the anion Ar ?
8.6
VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
Ionization energy and electron affinity help chemists understand the types of reactions
that elements undergo and the nature of the elements compounds. On a conceptual
level these two measures are related in a simple way: Ionization energy indexes the attraction of an atom for its own electrons, whereas electron affinity expresses the attraction of an atom for an additional electron from some other source. Together they
give us insight into the general attraction of an atom for electrons. With these concepts
we can survey the chemical behavior of the elements systematically, paying particular
attention to the relationship between chemical properties and electron configuration.
We have seen that the metallic character of the elements decreases from left to
right across a period and increases from top to bottom within a group. On the basis of
these trends and the knowledge that metals usually have low ionization energies while
nonmetals usually have high electron affinities, we can frequently predict the outcome
of a reaction involving some of these elements.
GENERAL TRENDS IN CHEMICAL PROPERTIES
1A
2A
3A
4A
Li
Be
B
C
Na
Mg
Al
Si
FIGURE 8.13 Diagonal relationships in the periodic table.
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Before we study the elements in individual groups, let us look at some overall trends.
We said have that elements in the same group resemble one another in chemical behavior because they have similar outer electron configurations. This statement, although
correct in the general sense, must be applied with caution. Chemists have long known
that the first member of each group (the element in the second period from lithium to
fluorine) differs from the rest of the members of the same group. Lithium, for example, exhibits many, but not all, of the properties characteristic of the alkali metals.
Similarly, beryllium is a somewhat atypical member of Group 2A, and so on. The difference can be attributed to the unusually small size of the first element in each group
(see Figure 8.5).
Another trend in the chemical behavior of the representative elements is the diagonal relationship. Diagonal relationships are similarities between pairs of elements
in different groups and periods of the periodic table. Specifically, the first three members of the second period (Li, Be, and B) exhibit many similarities to those elements
located diagonally below them in the periodic table (Figure 8.13). The reason for this
phenomenon is the closeness of the charge densities of their cations. (Charge density
is the charge of an ion divided by its volume.) Elements with comparable charge densities react similarly with anions and therefore form the same type of compounds. Thus
the chemistry of lithium resembles that of magnesium in some ways; the same holds
for beryllium and aluminum and for boron and silicon. Each of these pairs is said to
exhibit a diagonal relationship. We will see a number of examples of this relationship
later.
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
309
Bear in mind that a comparison of the properties of elements in the same group
is most valid if we are dealing with elements of the same type with respect to their
metallic character. This guideline applies to the elements in Groups 1A and 2A, which
are all metals, and to the elements in Groups 7A and 8A, which are all nonmetals. In
Groups 3A through 6A, where the elements change either from nonmetals to metals
or from nonmetals to metalloids, it is natural to expect greater variation in chemical
properties even though the members of the same group have similar outer electron configurations.
Now let us take a closer look at the chemical properties of the representative elements and the noble gases. (We will consider the chemistry of the transition metals
in Chapter 22.)
Hydrogen (1s1)
There is no totally suitable position for hydrogen in the periodic table. Traditionally
hydrogen is shown in Group 1A, but it really could be a class by itself. Like the alkali
metals, it has a single s valence electron and forms a unipositive ion (H ), which is
hydrated in solution. On the other hand, hydrogen also forms the hydride ion (H ) in
ionic compounds such as NaH and CaH2. In this respect, hydrogen resembles the halogens, all of which form uninegative ions (F , Cl , Br , and I ) in ionic compounds.
Ionic hydrides react with water to produce hydrogen gas and the corresponding metal
hydroxides:
2NaH(s)
2H2O(l) 88n 2NaOH(aq)
CaH2(s)
2H2O(l) 88n Ca(OH)2(s)
H2(g)
2H2(g)
Of course, the most important compound of hydrogen is water, which forms when hydrogen burns in air:
2H2(g)
Li
Na
K
Group 1A Elements (ns1, n
O2(g) 88n 2H2O(l)
2)
Figure 8.14 shows the Group 1A elements, the alkali metals. All of these elements have
low ionization energies and therefore a great tendency to lose the single valence electron. In fact, in the vast majority of their compounds they are unipositive ions. These
metals are so reactive that they are never found in the pure state in nature. They react
with water to produce hydrogen gas and the corresponding metal hydroxide:
Rb
2M(s)
Cs
2H2O(l) 88n 2MOH(aq)
H2(g)
where M denotes an alkali metal. When exposed to air, they gradually lose their shiny
appearance as they combine with oxygen gas to form oxides. Lithium forms lithium
oxide (containing the O2 ion):
4Li(s)
O2(g) 88n 2Li2O(s)
The other alkali metals all form oxides and peroxides (containing the O2 ion). For
2
example,
2Na(s)
O2(g) 88n Na2O2(s)
Potassium, rubidium, and cesium also form superoxides (containing the O2 ion):
K(s)
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O2(g) 88n KO2(s)
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.14 The Group 1A
elements: the alkali metals.
Francium (not shown) is radioactive.
Lithium (Li)
Rubidium (Rb)
Potassium (K)
Sodium (Na)
Cesium (Cs)
The reason that different types of oxides are formed when alkali metals react with oxygen has to do with the stability of the oxides in the solid state. Since these oxides are
all ionic compounds, their stability depends on how strongly the cations and anions attract one another. Lithium tends to form predominantly lithium oxide because this compound is more stable than lithium peroxide. The formation of other alkali metal oxides
can be explained similarly.
Be
Mg
Ca
Sr
Ba
Group 2A Elements (ns2, n
2)
Figure 8.15 shows the Group 2A elements. As a group, the alkaline earth metals are
somewhat less reactive than the alkali metals. Both the first and the second ionization
energies decrease from beryllium to barium. Thus the tendency is to form M2 ions
(where M denotes an alkaline earth metal atom), and hence the metallic character increases from top to bottom. Most beryllium compounds (BeH2 and beryllium halides,
such as BeCl2) and some magnesium compounds (MgH2, for example) are molecular
rather than ionic in nature.
The reactivities of alkaline earth metals with water vary quite markedly. Beryllium
does not react with water; magnesium reacts slowly with steam; calcium, strontium,
and barium are reactive enough to attack cold water:
Ba(s)
2H2O(l) 88n Ba(OH)2(aq)
H2(g)
The reactivities of the alkaline earth metals toward oxygen also increase from Be to
Ba. Beryllium and magnesium form oxides (BeO and MgO) only at elevated temperatures, whereas CaO, SrO, and BaO form at room temperature.
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
Beryllium (Be)
Magnesium (Mg)
Strontium (Sr)
Barium (Ba)
FIGURE 8.15 The Group 2A
elements: the alkaline earth
metals.
311
Calcium (Ca)
Radium (Ra)
Magnesium reacts with acids in aqueous solution, liberating hydrogen gas:
Mg(s)
2H (aq) 88n Mg2 (aq)
H2(g)
Calcium, strontium, and barium also react with aqueous acid solutions to produce hydrogen gas. However, because these metals also attack water, two different reactions
will occur simultaneously.
The chemical properties of calcium and strontium provide an interesting example
of periodic group similarity. Strontium-90, a radioactive isotope, is a major product of
an atomic bomb explosion. If an atomic bomb is exploded in the atmosphere, the strontium-90 formed will eventually settle on land and water, and it will reach our bodies
via a relatively short food chain. For example, if cows eat contaminated grass and drink
contaminated water, they will pass along strontium-90 in their milk. Because calcium
and strontium are chemically similar, Sr2 ions can replace Ca2 ions in our bones.
Constant exposure of the body to the high-energy radiation emitted by the strontium90 isotopes can lead to anemia, leukemia, and other chronic illnesses.
B
Al
Ga
Group 3A Elements (ns2np1, n
The first member of Group 3A, boron, is a metalloid; the rest are metals (Figure 8.16).
Boron does not form binary ionic compounds and is unreactive toward oxygen gas and
water. The next element, aluminum, readily forms aluminum oxide when exposed to
air:
4Al(s)
In
Tl
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2)
3O2(g) 88n 2Al2O3(s)
Aluminum that has a protective coating of aluminum oxide is less reactive than elemental aluminum. Aluminum forms only tripositive ions. It reacts with hydrochloric
acid as follows:
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.16
elements.
The Group 3A
Boron (B)
Aluminum (Al)
Gallium (Ga)
Indium (In)
2Al(s)
6H (aq) 88n 2Al3 (aq)
3H2(g)
The other Group 3A metallic elements form both unipositive and tripositive ions.
Moving down the group, we find that the unipositive ion becomes more stable than the
tripositive ion.
The metallic elements in Group 3A also form many molecular compounds. For
example, aluminum reacts with hydrogen to form AlH3, which resembles BeH2 in its
properties. (Here is an example of the diagonal relationship.) Thus, from left to right
across the periodic table, we are seeing a gradual shift from metallic to nonmetallic
character in the representative elements.
C
Si
Ge
Group 4A Elements (ns2np2, n
The first member of group 4A, carbon, is a nonmetal, and the next two members, silicon and germanium, are metalloids (Figure 8.17). These elements do not form ionic
compounds. The metallic elements of this group, tin and lead, do not react with water,
but they do react with acids (hydrochloric acid, for example) to liberate hydrogen gas:
Sn(s)
Sn
Pb(s)
Pb
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2)
2H (aq) 88n Sn2 (aq)
2
2H (aq) 88n Pb (aq)
H2(g)
H2(g)
The Group 4A elements form compounds in both the 2 and 4 oxidation states.
For carbon and silicon, the 4 oxidation state is the more stable one. For example, CO2
is more stable than CO, and SiO2 is a stable compound, but SiO does not exist under
normal conditions. As we move down the group, however, the trend in stability is
reversed. In tin compounds the 4 oxidation state is only slightly more stable than the
2 oxidation state. In lead compounds the 2 oxidation state is unquestionably the
more stable one. The outer electron configuration of lead is 6s26p2, and lead tends to
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
Carbon (graphite)
Carbon (diamond)
Germanium (Ge)
FIGURE 8.17 The Group 4A
elements.
N
P
As
Sb
Bi
Silicon (Si)
Tin (Sn)
Lead (Pb)
lose only the 6p electrons (to form Pb2 ) rather than both the 6p and 6s electrons (to
form Pb4 ).
Group 5A Elements (ns2np3, n
P4O10(s)
S
Se
Te
Po
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2)
In Group 5A, nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a metal (Figure 8.18). Thus we expect a greater variation in
properties within the group.
Elemental nitrogen is a diatomic gas (N2). It forms a number of oxides (NO, N2O,
NO2, N2O4, and N2O5), of which only N2O5 is a solid; the others are gases. Nitrogen
has a tendency to accept three electrons to form the nitride ion, N3 (thus achieving the
electron configuration 2s22p6, which is isoelectronic with neon). Most metallic nitrides
(Li3N and Mg3N2, for example) are ionic compounds. Phosphorus exists as P4 molecules. It forms two solid oxides with the formulas P4O6 and P4O10. The important
oxoacids HNO3 and H3PO4 are formed when the following oxides react with water:
N2O5(s)
O
313
H2O(l) 88n 2HNO3(aq)
6H2O(l) 88n 4H3PO4(aq)
Arsenic, antimony, and bismuth have extensive three-dimensional structures. Bismuth
is a far less reactive metal than those in the preceding groups.
Group 6A Elements (ns2np4, n
2)
The first three members of Group 6A (oxygen, sulfur, and selenium) are nonmetals,
and the last two (tellurium and polonium) are metalloids (Figure 8.19). Oxygen is a
diatomic gas; elemental sulfur and selenium have the molecular formulas S8 and Se8,
respectively; tellurium and polonium have more extensive three-dimensional structures.
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
FIGURE 8.18 The Group 5A
elements. Molecular nitrogen is a
colorless, odorless gas.
in
e
bl
a
ail
Av
ot
N
on
rsi
e
tV
x
Te
e-
Nitrogen (N2)
Arsenic (As)
Antimony (Sb)
Sulfur (S8)
Selenium (Se8)
White and red phosphorous (P)
Bismuth (Bi)
Tellurium (Te)
FIGURE 8.19 The Group 6A elements sulfur, selenium, and tellurium. Molecular oxygen is a
colorless, odorless gas. Polonium (not shown) is radioactive.
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
315
FIGURE 8.20 The Group 7A
elements chlorine, bromine, and
iodine. Fluorine is a greenishyellow gas that attacks ordinary glassware. Astatine is
radioactive.
(Polonium is a radioactive element that is difficult to study in the laboratory.) Oxygen
has a tendency to accept two electrons to form the oxide ion (O2 ) in many ionic compounds. Sulfur, selenium, and tellurium also form dinegative anions (S2 , Se2 , and
Te2 ). The elements in this group (especially oxygen) form a large number of molecular compounds with nonmetals. The important compounds of sulfur are SO2, SO3,
and H2S. Sulfuric acid is formed when sulfur trioxide reacts with water:
SO3(g)
F
Cl
Br
Group 7A Elements (ns2np5, n
H2O(l ) 88n H2SO4(aq)
2)
All the halogens are nonmetals with the general formula X2, where X denotes a halogen element (Figure 8.20). Because of their great reactivity, the halogens are never
found in the elemental form in nature. (The last member of Group 7A, astatine, is a
radioactive element. Little is known about its properties.) Fluorine is so reactive that
it attacks water to generate oxygen:
I
2F2(g)
At
2H2O(l) 88n 4HF(aq)
O2(g)
Actually the reaction between molecular fluorine and water is quite complex; the products formed depend on reaction conditions. The reaction shown above is one of several possible changes.
The halogens have high ionization energies and high electron affinities. Anions
derived from the halogens (F , Cl , Br , and I ) are called halides. They are isoelectronic with the noble gases immediately to their right in the periodic table. For example, F is isoelectronic with Ne, Cl with Ar, and so on. The vast majority of the
alkali metal halides and alkaline earth metal halides are ionic compounds. The halogens also form many molecular compounds among themselves (such as ICl and BrF3)
and with nonmetallic elements in other groups (such as NF3, PCl5, and SF6). The halogens react with hydrogen to form hydrogen halides:
H2(g)
X2(g) 88n 2HX(g)
When this reaction involves fluorine, it is explosive, but it becomes less and less violent as we substitute chlorine, bromine, and iodine. The hydrogen halides dissolve in
water to form hydrohalic acids. Hydrofluoric acid (HF) is a weak acid (that is, it is a
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
Helium (He)
Neon (Ne)
FIGURE 8.21 All noble gases
are colorless and odorless. These
pictures show the colors emitted
by the gases from a discharge
tube.
He
Ne
Ar
Kr
Xe
Argon (Ar)
Krypton (Kr)
Xenon (Xe)
weak electrolyte), but the other hydrohalic acids (HCl, HBr, and HI) are all strong acids
(strong electrolytes).
Group 8A Elements (ns2np6, n
2)
All noble gases exist as monatomic species (Figure 8.21). Their atoms have completely
filled outer ns and np subshells, which give them great stability. (Helium is 1s2.) The
Group 8A ionization energies are among the highest of all elements, and these gases
have no tendency to accept extra electrons. For years these elements were called inert
gases, and rightly so. Until 1963 no one had been able to prepare a compound containing any of these elements. The British chemist Neil Bartlett shattered chemists
long-held views of these elements when he exposed xenon to platinum hexafluoride,
a strong oxidizing agent, and brought about the following reaction (Figure 8.22):
Xe(g)
PtF6(g) 88n XePtF6(s)
Since then, a number of xenon compounds (XeF4, XeO3, XeO4, XeOF4) and a few
krypton compounds (KrF2, for example) have been prepared (Figure 8.23). Despite the
immense interest in the chemistry of the noble gases, however, their compounds do not
have any commercial applications, and they are not involved in natural biological
processes. No compounds of helium, neon, and argon are known.
COMPARISON OF GROUP 1A AND GROUP 1B ELEMENTS
When we compare the Group 1A elements (alkali metals) and the Group 1B elements
(copper, silver, and gold), we arrive at an interesting conclusion. Although the metals
in these two groups have similar outer electron configurations, with one electron in the
outermost s orbital, their chemical properties are quite different.
Neil Bartlett (1932 ). English chemist. Bartletts work is mainly in the preparation and study of compounds with unusual oxidation states and in solid-state chemistry.
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
317
FIGURE 8.22 (a) Xenon gas
(colorless) and PtF6 (red gas) separated from each other. (b) When
the two gases are allowed to
mix, a yellow-orange solid compound is formed.
(a)
FIGURE 8.23 Crystals of
xenon tetrafluoride (XeF4).
(b)
The first ionization energies of Cu, Ag, and Au are 745 kJ/mol, 731 kJ/mol, and
890 kJ/mol, respectively. Since these values are considerably larger than those of the
alkali metals (see Table 8.3), the Group 1B elements are much less reactive. The higher
ionization energies of the Group 1B elements result from incomplete shielding of the
nucleus by the inner d electrons (compared with the more effective shielding of the
completely filled noble gas cores). Consequently the outer s electrons of these elements
are more strongly attracted by the nucleus. In fact, copper, silver, and gold are so unreactive that they are usually found in the uncombined state in nature. The inertness
and rarity of these metals make them valuable in the manufacture of coins and in jewelry. For this reason, these metals are also called coinage metals. The difference in
chemical properties between the Group 2A elements (the alkaline earth metals) and the
Group 2B metals (zinc, cadmium, and mercury) can be explained in a similar way.
PROPERTIES OF OXIDES ACROSS A PERIOD
One way to compare the properties of the representative elements across a period is to
examine the properties of a series of similar compounds. Since oxygen combines with
almost all elements, we will compare the properties of oxides of the third-period elements to see how metals differ from metalloids and nonmetals. Some elements in the
third period (P, S, and Cl) form several types of oxides, but for simplicity we will consider only those oxides in which the elements have the highest oxidation number. Table
8.5 lists a few general characteristics of these oxides. We observed earlier that oxygen
Some Properties of Oxides of the Third-Period Elements
Na2O
MgO
Al2O3
SiO2
P4O10
SO3
Cl2O7
m8888
m888888
ABLE 8.5
m8888888
Type of compound
Structure
Melting point (C)
Boiling point (C)
Acid-base nature
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2800
3600
Basic
2045
2980
Amphoteric
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molecular units
16.8
91.5
44.8
82
m8888888 Acidic
1610
2230
580
?
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1275
?
Basic
m88
Back
m8888888 Ionic
m888888 Molecular
m88 Extensive three-dimensional
m888 Discrete
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
has a tendency to form the oxide ion. This tendency is greatly favored when oxygen
combines with metals that have low ionization energies, namely, those in Groups 1A
and 2A, plus aluminum. Thus Na2O, MgO, and Al2O3 are ionic compounds, as indicated by their high melting points and boiling points. They have extensive threedimensional structures in which each cation is surrounded by a specific number of anions, and vice versa. As the ionization energies of the elements increase from left to
right, so does the molecular nature of the oxides that are formed. Silicon is a metalloid; its oxide (SiO2) also has a huge three-dimensional network, although no ions are
present. The oxides of phosphorus, sulfur, and chlorine are molecular compounds composed of small discrete units. The weak attractions among these molecules result in
relatively low melting points and boiling points.
Most oxides can be classified as acidic or basic depending on whether they produce acids or bases when dissolved in water or react as acids or bases in certain
processes. Some oxides are amphoteric, which means that they display both acidic and
basic properties. The first two oxides of the third period, Na2O and MgO, are basic
oxides. For example, Na2O reacts with water to form the base sodium hydroxide:
Na2O(s)
H2O(l) 88n 2NaOH(aq)
Magnesium oxide is quite insoluble; it does not react with water to any appreciable extent. However, it does react with acids in a manner that resembles an acid-base reaction:
MgO(s)
2HCl(aq) 88n MgCl2(aq)
H2O(l)
Note that the products of this reaction are a salt (MgCl2) and water, the usual products
of an acid-base neutralization.
Aluminum oxide is even less soluble than magnesium oxide; it too does not react with water. However, it shows basic properties by reacting with acids:
Al2O3(s)
6HCl(aq) 88n 2AlCl3(aq)
3H2O(l)
It also exhibits acidic properties by reacting with bases:
Al2O3(s)
Note that this acid-base neutralization produces a salt but no water.
2NaOH(aq)
3H2O(l) 88n 2NaAl(OH)4(aq)
Thus Al2O3 is classified as an amphoteric oxide because it has properties of both acids
and bases. Other amphoteric oxides are ZnO, BeO, and Bi2O3.
Silicon dioxide is insoluble and does not react with water. It has acidic properties, however, because it reacts with very concentrated bases:
SiO2(s)
2NaOH(aq) 88n Na2SiO3(aq)
H2O(l)
For this reason, concentrated aqueous, strong bases such as NaOH(aq) should not be
stored in Pyrex glassware, which is made of SiO2.
The remaining third-period oxides are acidic. They react with water to form phosphoric acid (H3PO4), sulfuric acid (H2SO4), and perchloric acid (HClO4):
P4O10(s)
SO3(g)
Cl2O7(l )
6H2O(l) 88n 4H3PO4(aq)
H2O(l) 88n H2SO4(aq)
H2O(l) 88n 2HClO4(aq)
This brief examination of oxides of the third-period elements shows that as the
metallic character of the elements decreases from left to right across the period, their
oxides change from basic to amphoteric to acidic. Metallic oxides are usually basic,
and most oxides of nonmetals are acidic. The intermediate properties of the oxides (as
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VARIATION IN CHEMICAL PROPERTIES OF THE REPRESENTATIVE ELEMENTS
319
shown by the amphoteric oxides) are exhibited by elements whose positions are intermediate within the period. Note also that since the metallic character of the elements
increases from top to bottom within a group of representative elements, we would expect oxides of elements with higher atomic numbers to be more basic than the lighter
elements. This is indeed the case.
EXAMPLE 8.6
Classify the following oxides as acidic, basic, or amphoteric: (a) Rb2O, (b) BeO,
(c) As2O5.
(a) Since rubidium is an alkali metal, we would expect Rb2O to be a basic oxide. This is indeed true, as shown by rubidium oxides reaction with water to
form rubidium hydroxide:
Answer
Rb2O(s)
H2O(l) 88n 2RbOH(aq)
(b) Beryllium is an alkaline earth metal. However, because it is the first member of
Group 2A, we expect that it may differ somewhat from the other members of the
group. Furthermore, beryllium and aluminum exhibit a diagonal relationship, so that
BeO may resemble Al2O3 in properties. It turns out that BeO, like Al2O3, is an amphoteric oxide, as shown by its reactions with acids and bases:
BeO(s)
2H (aq)
3H2O(l) 88n Be(H2O)2 (aq)
4
BeO(s)
2OH (aq)
H2O(l) 88n Be(OH)2 (aq)
4
(c) Since arsenic is a nonmetal, we expect As2O5 to be an acidic oxide. This prediction is correct, as shown by the formation of arsenic acid when As2O5 reacts with
water:
As2O5(s)
Similar problem: 8.70.
3H2O(l) 88n 2H3AsO4(aq)
PRACTICE EXERCISE
Classify the following oxides as acidic, basic, or amphoteric: (a) ZnO, (b) P4O10,
(c) CaO.
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Discovery of the Noble Gases
In the late 1800s John William Strutt, Third Baron of
Rayleigh, who was a professor of physics at the
Cavendish Laboratory in Cambridge, England, accurately determined the atomic masses of a number of
elements, but he obtained a puzzling result with nitrogen. One of his methods of preparing nitrogen was
by the thermal decomposition of ammonia:
2NH3(g) 88n N2(g)
3H2(g)
it oxygen, carbon dioxide, and water vapor.
Invariably, the nitrogen from air was a little denser (by
about 0.5%) than the nitrogen from ammonia.
Lord Rayleighs work caught the attention of Sir
William Ramsay, a professor of chemistry at the
University College, London. In 1898 Ramsay passed
nitrogen, which he had obtained from air by
Rayleighs procedure, over red-hot magnesium to convert it to magnesium nitride:
Another method was to start with air and remove from
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3Mg(s)
N2(g) 88n Mg3N2(s)
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
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After all of the nitrogen had reacted with magnesium,
Ramsey was left with an unknown gas that would not
combine with anything.
With the help of Sir William Crookes, the inventor of the discharge tube, Ramsay and Lord Rayleigh
found that the emission spectrum of the gas did not
match any of the known elements. The gas was a new
element! They determined its atomic mass to be
39.95 amu and called it argon, which means the
lazy one in Greek.
Once argon had been discovered, other noble
gases were quickly identified. Also in 1898 Ramsay
isolated helium from uranium ores (see Chemistry in
Action essay on p. 255). From the atomic masses of
helium and argon, their lack of chemical reactivity,
and what was then known about the periodic table,
Ramsay was convinced that there were other unreactive gases and that they were all members of one periodic group. He and his student Morris Travers set
out to find the unknown gases. They used a refrigeration machine to first produce liquid air. Applying a
technique called fractional distillation, they then allowed the liquid air to warm up gradually and collected components that boiled off at different temperatures. In this manner they analyzed and identified
three new elements neon, krypton, and xenon in
only three months. Three new elements in three months
is a record that may never be broken!
The discovery of the noble gases helped to complete the periodic table. Their atomic masses suggested that these elements should be placed to the
right of the halogens. The apparent discrepancy with
the position of argon was resolved by Moseley, as discussed in the chapter.
Finally, the last member of the noble gases, radon,
was discovered by the German chemist Frederick Dorn
SUMMARY OF FACTS
AND CONCEPTS
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Sir William Ramsay (18521916).
in 1900. A radioactive element and the heaviest elemental gas known, radons discovery not only completed the Group 8A elements, but also advanced our
understanding about the nature of radioactive decay
and transmutation of elements.
Lord Rayleigh and Ramsay both won Nobel Prizes
in 1904 for the discovery of argon. Lord Rayleigh received the prize in physics and Ramsays award was
in chemistry.
1. Nineteenth-century chemists developed the periodic table by arranging elements in the increasing order of their atomic masses. Discrepancies in early versions of the periodic table
were resolved by arranging the elements in order of their atomic numbers.
2. Electron configuration determines the properties of an element. The modern periodic table
classifies the elements according to their atomic numbers, and thus also by their electron
configurations. The configuration of the valence electrons directly affects the properties of
the atoms of the representative elements.
3. Periodic variations in the physical properties of the elements reflect differences in atomic
structure. The metallic character of elements decreases across a period from metals through
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QUESTIONS AND PROBLEMS
321
the metalloids to nonmetals and increases from top to bottom within a particular group of
representative elements.
4. Atomic radius varies periodically with the arrangement of the elements in the periodic
table. It decreases from left to right and increases from top to bottom.
5. Ionization energy is a measure of the tendency of an atom to resist the loss of an electron.
The higher the ionization energy, the stronger the attraction between the nucleus and an electron. Electron affinity is a measure of the tendency of an atom to gain an electron. The more
positive the electron affinity, the greater the tendency for the atom to gain an electron. Metals
usually have low ionization energies, and nonmetals usually have high electron affinities.
6. Noble gases are very stable because their outer ns and np subshells are completely filled. The
metals among the representative elements (in Groups 1A, 2A, and 3A) tend to lose electrons
until their cations become isoelectronic with the noble gases that precede them in the periodic
table. The nonmetals in Groups 5A, 6A, and 7A tend to accept electrons until their anions
become isoelectronic with the noble gases that follow them in the periodic table.
KEY WORDS
Amphoteric oxide, p. 318
Atomic radius, p. 295
Diagonal relationship, p. 308
Electron affinity, p. 305
Ionic radius, p. 297
Ionization energy, p. 301
Isoelectronic, p. 294
Valence electrons, p. 291
Noble gases, p. 290
Representative elements, p. 290
QUESTIONS AND PROBLEMS
DEVELOPMENT OF THE PERIODIC TABLE
Review Questions
8.1 Briefly describe the significance of Mendeleevs periodic table.
8.2 What is Moseleys contribution to the modern periodic table?
8.3 Describe the general layout of a modern periodic
table.
8.4 What is the most important relationship among elements in the same group in the periodic table?
PERIODIC CLASSIFICATION OF THE ELEMENTS
Review Questions
8.5 Which of the following elements are metals, nonmetals, or metalloids? As, Xe, Fe, Li, B, Cl, Ba, P, I,
Si.
8.6 Compare the physical and chemical properties of metals and nonmetals.
8.7 Draw a rough sketch of a periodic table (no details
are required). Indicate regions where metals, nonmetals, and metalloids are located.
8.8 What is a representative element? Give names and
symbols of four representative elements.
8.9 Without referring to a periodic table, write the name
and give the symbol for an element in each of the following groups: 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A,
transition metals.
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8.10 Indicate whether the following elements exist as
atomic species, molecular species, or extensive threedimensional structures in their most stable states at
25C and 1 atm and write the molecular or empirical
formula for each one: phosphorus, iodine, magnesium, neon, arsenic, sulfur, boron, selenium, and oxygen.
8.11 You are given a dark shiny solid and asked to determine whether it is iodine or a metallic element.
Suggest a nondestructive test that would allow you to
arrive at the correct answer.
8.12 What are valence electrons? For representative elements, the number of valence electrons of an element
is equal to its group number. Show that this is true
for the following elements: Al, Sr, K, Br, P, S, C.
8.13 Write the outer electron configurations for the (a) alkali metals, (b) alkaline earth metals, (c) halogens,
(d) noble gases.
8.14 Use the first-row transition metals (Sc to Cu) as an
example to illustrate the characteristics of the electron configurations of transition metals.
8.15 How does the electron configuration of ions derived
from representative elements give them stability?
8.16 What do we mean when we say that two ions or an
atom and an ion are isoelectronic?
8.17 What is wrong with the statement The atoms of element X are isoelectronic with the atoms of element
Y?
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
8.18 Give three examples of first-row transition metal (Sc
to Cu) ions whose electron configurations are represented by the argon core.
8.29
Problems
8.19 In the periodic table, the element hydrogen is sometimes grouped with the alkali metals (as in this book)
and sometimes with the halogens. Explain why hydrogen can resemble the Group 1A and the Group 7A
elements.
8.20 A neutral atom of a certain element has 17 electrons.
Without consulting a periodic table, (a) write the
ground-state electron configuration of the element,
(b) classify the element, (c) determine whether the
atoms of this element are diamagnetic or paramagnetic.
8.21 Group the following electron configurations in pairs
that would represent similar chemical properties of
their atoms:
(a) 1s22s22p63s2
(b) 1s22s22p3
(c) 1s22s22p63s23p64s23d104p6
(d) 1s22s2
(e) 1s22s22p6
(f) 1s22s22p63s23p3
8.22 Group the following electron configurations in pairs
that would represent similar chemical properties of
their atoms:
(a) 1s22s22p5
(b) 1s22s1
(c) 1s22s22p6
(d) 1s22s22p63s23p5
(e) 1s22s22p63s23p64s1
(f) 1s22s22p63s23p64s23d104p6
8.23 Without referring to a periodic table, write the electron configuration of elements with the following
atomic numbers: (a) 9, (b) 20, (c) 26, (d) 33. Classify
the elements.
8.24 Specify the group of the periodic table in which each
of the following elements is found: (a) [Ne]3s1,
(b) [Ne]3s23p3, (c) [Ne]3s23p6, (d) [Ar]4s23d8.
8.25 A M2 ion derived from a metal in the first transition metal series has four electrons in the 3d subshell.
What element might M be?
8.26 A metal ion with a net 3 charge has five electrons
in the 3d subshell. Identify the metal.
8.27 Write the ground-state electron configurations of the
following ions: (a) Li , (b) H , (c) N3 , (d) F ,
(e) S2 , (f) Al3 , (g) Se2 , (h) Br , (i) Rb , (j) Sr2 ,
(k) Sn2 , (l) Te2 , (m) Ba2 , (n) Pb2 , (o) In3 ,
(p) Tl , (q) Tl3 .
8.28 Write the ground-state electron configurations of the
following ions, which play important roles in bio-
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8.30
8.31
8.32
chemical processes in our bodies: (a) Na , (b) Mg2 ,
(c) Cl , (d) K , (e) Ca2 , (f) Fe2 , (g) Cu2 ,
(h) Zn2 .
Write the ground-state electron configurations of the
following transition metal ions: (a) Sc3 , (b) Ti4 ,
(c) V5 , (d) Cr3 , (e) Mn2 , (f) Fe2 , (g) Fe3 ,
(h) Co2 , (i) Ni2 , (j) Cu , (k) Cu2 , (l) Ag ,
(m) Au , (n) Au3 , (o) Pt2 .
Name the ions with 3 charges that have the following electron configurations: (a) [Ar]3d3, (b) [Ar],
(c) [Kr]4d6, (d) [Xe]4f145d6.
Which of the following species are isoelectronic with
each other? C, Cl , Mn2 , B , Ar, Zn, Fe3 , Ge2 .
Group the species that are isoelectronic: Be2 , F ,
Fe2 , N3 , He, S2 , Co3 , Ar.
PERIODIC VARIATION IN PHYSICAL PROPERTIES
Review Questions
8.33 Define atomic radius. Does the size of an atom have
a precise meaning?
8.34 How does atomic radius change (a) from left to right
across a period and (b) from top to bottom in a group?
8.35 Define ionic radius. How does the size of an atom
change when it is converted to (a) an anion and (b) a
cation?
8.36 Explain why, for isoelectronic ions, the anions are
larger than the cations.
Problems
8.37 On the basis of their positions in the periodic table,
select the atom with the larger atomic radius in each
of the following pairs: (a) Na, Cs; (b) Be, Ba;
(c) N, Sb; (d) F, Br; (e) Ne, Xe.
8.38 Arrange the following atoms in order of decreasing
atomic radius: Na, Al, P, Cl, Mg.
8.39 Which is the largest atom in Group 4A?
8.40 Which is the smallest atom in Group 7A?
8.41 Why is the radius of the lithium atom considerably
larger than the radius of the hydrogen atom?
8.42 Use the second period of the periodic table as an example to show that the size of atoms decreases as we
move from left to right. Explain the trend.
8.43 Indicate which one of the two species in each of the
following pairs is smaller: (a) Cl or Cl ; (b) Na or
Na ; (c) O2 or S2 ; (d) Mg2 or Al3 ; (e) Au or
Au3 .
8.44 List the following ions in order of increasing ionic
radius: N3 , Na , F , Mg2 , O2 .
8.45 Explain which of the following cations is larger, and
why: Cu or Cu2 .
8.46 Explain which of the following anions is larger, and
why: Se2 or Te2 .
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8.47 Give the physical states (gas, liquid, or solid) of the
representative elements in the fourth period (K, Ca,
Ga, Ge, As, Se, Br) at 1 atm and 25C.
8.48 The boiling points of neon and krypton are 245.9C
and 152.9C, respectively. Using these data, estimate the boiling point of argon.
IONIZATION ENERGY
Review Questions
8.49 Define ionization energy. Ionization energy measurements are usually made when atoms are in the
gaseous state. Why? Why is the second ionization energy always greater than the first ionization energy
for any element?
8.50 Sketch the outline of the periodic table and show
group and period trends in the first ionization energy
of the elements. What types of elements have the
highest ionization energies and what types the lowest ionization energies?
Problems
8.51 Use the third period of the periodic table as an example to illustrate the change in first ionization energies of the elements as we move from left to right.
Explain the trend.
8.52 In general, ionization energy increases from left to
right across a given period. Aluminum, however, has
a lower ionization energy than magnesium. Explain.
8.53 The first and second ionization energies of K are
419 kJ/mol and 3052 kJ/mol, and those of Ca are
590 kJ/mol and 1145 kJ/mol, respectively. Compare
their values and comment on the differences.
8.54 Two atoms have the electron configurations 1s22s22p6
and 1s22s22p63s1. The first ionization energy of one
is 2080 kJ/mol, and that of the other is 496 kJ/mol.
Match each ionization energy with one of the given
electron configurations. Justify your choice.
8.55 A hydrogenlike ion is an ion containing only one electron. The energies of the electron in a hydrogenlike
ion are given by
En
(2.18
10
18
J)Z 2
1
n2
where n is the principal quantum number and Z is the
atomic number of the element. Calculate the ionization energy (in kJ/mol) of the He ion.
8.56 Plasma is a state of matter consisting of positive
gaseous ions and electrons. In the plasma state, a mercury atom could be stripped of its 80 electrons and
therefore would exist as Hg80 . Use the equation in
Problem 8.55 to calculate the energy required for the
last ionization step, that is,
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Hg79 (g) 88n Hg80 (g)
323
e
ELECTRON AFFINITY
Review Questions
8.57 (a) Define electron affinity. (b) Electron affinity measurements are made with gaseous atoms. Why?
(c) Ionization energy is always a positive quantity,
whereas electron affinity may be either positive or
negative. Explain.
8.58 Explain the trends in electron affinity from aluminum
to chlorine (see Table 8.4).
Problems
8.59 Arrange the elements in each of the following groups
in increasing order of the most positive electron affinity: (a) Li, Na, K, (b) F, Cl, Br, I.
8.60 Specify which of the following elements you would
expect to have the greatest electron affinity: He, K,
Co, S, Cl.
8.61 Considering their electron affinities, do you think it
is possible for the alkali metals to form an anion like
M , where M represents an alkali metal?
8.62 Explain why alkali metals have a greater affinity for
electrons than alkaline earth metals.
VARIATION IN CHEMICAL PROPERTIES OF THE
REPRESENTATIVE ELEMENTS
Review Questions
8.63 What is meant by the diagonal relationship? Name
two pairs of elements that show this relationship.
8.64 Which elements are more likely to form acidic oxides? Basic oxides? Amphoteric oxides?
Problems
8.65 Use the alkali metals and alkaline earth metals as examples to show how we can predict the chemical
properties of elements simply from their electron configurations.
8.66 Based on your knowledge of the chemistry of the alkali metals, predict some of the chemical properties
of francium, the last member of the group.
8.67 As a group, the noble bases are very stable chemically (only Kr and Xe are known to form compounds).
Why?
8.68 Why are Group 1B elements more stable than Group
1A elements even though they seem to have the same
outer electron configuration, ns1, where n is the principal quantum number of the outermost shell?
8.69 How do the chemical properties of oxides change
from left to right across a period? From top to bottom within a particular group?
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
8.70 Write balanced equations for the reactions between
each of the following oxides and water: (a) Li2O,
(b) CaO, (c) SO3.
8.71 Write formulas for and name the binary hydrogen
compounds of the second-period elements (Li to F).
Describe how the physical and chemical properties
of these compounds change from left to right across
the period.
8.72 Which oxide is more basic, MgO or BaO? Why?
8.73 State whether each of the following properties of the
representative elements generally increases or decreases (a) from left to right across a period and
(b) from top to bottom within a group: metallic character, atomic size, ionization energy, acidity of oxides.
8.74 With reference to the periodic table, name (a) a halogen element in the fourth period, (b) an element similar to phosphorus in chemical properties, (c) the most
reactive metal in the fifth period, (d) an element that
has an atomic number smaller than 20 and is similar
to strontium.
8.75 Write equations representing the following processes:
(a) The electron affinity of S .
(b) The third ionization energy of titanium.
(c) The electron affinity of Mg2 .
(d) The ionization energy of O2 .
8.76 Arrange the following isoelectronic species in order
of (a) increasing ionic radius and (b) increasing ionization energy: O2 , F , Na , Mg2 .
8.77 Write the empirical (or molecular) formulas of compounds that the elements in the third period (sodium
to chlorine) should form with (a) molecular oxygen
and (b) molecular chlorine. In each case indicate
whether you would expect the compound to be ionic
or molecular in character.
8.78 Element M is a shiny and highly reactive metal (melting point 63C), and element X is a highly reactive
nonmetal (melting point 7.2C). They react to form
a compound with the empirical formula MX, a colorless, brittle white solid that melts at 734C. When
dissolved in water or when in the molten state, the
substance conducts electricity. When chlorine gas is
bubbled through an aqueous solution containing MX,
a reddish-brown liquid appears and Cl ions are
formed. From these observations, identify M and X.
(You may need to consult a handbook of chemistry
for the melting-point values.)
8.79 Match each of the elements on the right with its description on the left:
(a) A dark-red liquid
Calcium (Ca)
(b) A colorless gas that burns
Gold (Au)
in oxygen gas
Hydrogen (H2)
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8.81
8.82
ADDITIONAL PROBLEMS
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8.83
8.84
8.85
8.86
8.87
8.88
8.89
8.90
8.91
8.92
8.93
(c) A reactive metal that attacks
Argon (Ar)
water
Bromine (Br2)
(d) A shiny metal that is used in
jewelry
(e) A totally inert gas
Arrange the following species in isoelectronic pairs:
O , Ar, S2 , Ne, Zn, Cs , N3 , As3 , N, Xe.
In which of the following are the species written in
decreasing order by size of radius? (a) Be, Mg, Ba,
(b) N3 , O2 , F , (c) Tl3 , Tl2 , Tl .
Which of the following properties show a clear periodic variation? (a) first ionization energy, (b) molar
mass of the elements, (c) number of isotopes of an
element, (d) atomic radius.
When carbon dioxide is bubbled through a clear calcium hydroxide solution, the solution appears milky.
Write an equation for the reaction and explain how
this reaction illustrates that CO2 is an acidic oxide.
You are given four substances: a fuming red liquid,
a dark metallic-looking solid, a pale-yellow gas, and
a yellow-green gas that attacks glass. You are told that
these substances are the first four members of Group
7A, the halogens. Name each one.
For each pair of elements listed below, give three
properties that show their chemical similarity:
(a) sodium and potassium and (b) chlorine and
bromine.
Name the element that forms compounds, under appropriate conditions, with every other element in the
periodic table except He, Ne, and Ar.
Explain why the first electron affinity of sulfur is
200 kJ/mol but the second electron affinity is
649 kJ/mol.
The H ion and the He atom have two 1s electrons
each. Which of the two species is larger? Explain.
Predict the products of the following oxides with water: Na2O, BaO, CO2, N2O5, P4O10, SO3. Write an
equation for each of the reactions. Specify whether
the oxides are acidic, basic, or amphoteric.
Write the formulas and names of the oxides of the
second-period elements (Li to N). Identify the oxides
as acidic, basic, or amphoteric.
State whether each of the following elements is a gas,
a liquid, or a solid under atmospheric conditions. Also
state whether it exists in the elemental form as atoms,
as molecules, or as a three-dimensional network: Mg,
Cl, Si, Kr, O, I, Hg, Br.
What factors account for the unique nature of hydrogen?
The air in a manned spacecraft or submarine needs
to be purified of exhaled carbon dioxide. Write equations for the reactions between carbon dioxide and
(a) lithium oxide (Li2O), (b) sodium peroxide
(Na2O2), and (c) potassium superoxide (KO2).
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8.94 The formula for calculating the energies of an electron in a hydrogenlike ion is given in Problem 8.55.
This equation cannot be applied to many-electron
atoms. One way to modify it for the more complex
), where Z is the
atoms is to replace Z with (Z
is a positive dimensionless
atomic number and
quantity called the shielding constant. Consider the
helium atom as an example. The physical significance
of is that it represents the extent of shielding that
the two 1s electrons exert on each other. Thus the
) is appropriately called the effective
quantity (Z
nuclear charge. Calculate the value of if the first
ionization energy of helium is 3.94 10 18 J per
atom. (Ignore the minus sign in the given equation in
your calculation.)
8.95 Why do noble gases have negative electron affinity
values?
8.96 The atomic radius of K is 216 pm and that of K is
133 pm. Calculate the percent decrease in volume that
occurs when K(g) is converted to K (g). [The volume of a sphere is ( 4 ) r3, where r is the radius of the
3
sphere.]
8.97 The atomic radius of F is 72 pm and that of F is
136 pm. Calculate the percent increase in volume that
occurs when F(g) is converted to F (g). (See Problem
8.96 for the volume of a sphere.)
8.98 A technique called photoelectron spectroscopy is
used to measure the ionization energy of atoms. A
sample is irradiated with UV light, and electrons are
ejected from the valence shell. The kinetic energies
of the ejected electrons are measured. Since the energy of the UV photon and the kinetic energy of the
ejected electron are known, we can write
h
IE
1
2
mu2
where is the frequency of the UV light, and m and
u are the mass and velocity of the electron, respectively. In one experiment the kinetic energy of the
ejected electron from potassium is found to be 5.34
10 19 J using a UV source of wavelength 162 nm.
Calculate the ionization energy of potassium. How
can you be sure that this ionization energy corresponds to the electron in the valence shell (that is, the
most loosely held electron)?
8.99 Referring to the Chemistry in Action essay on p. 319,
answer the following questions. (a) Why did it take
so long to discover the first noble gas (argon) on
Earth? (b) Once argon had been discovered, why did
it take relatively little time to discover the rest of the
noble gases? (c) Why was helium not isolated by the
fractional distillation of liquid air?
8.100 The energy needed for the following process is
1.96 104 kJ/mol:
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Li(g) 88n Li3 (g)
325
3e
If the first ionization energy of lithium is 520 kJ/mol,
calculate the second ionization energy of lithium, that
is, the energy required for the process
Li (g) 88n Li2 (g)
e
(Hint: You need the equation in Problem 8.55.)
8.101 An element X reacts with hydrogen gas at 200C to
form compound Y. When Y is heated to a higher temperature, it decomposes to the element X and hydrogen gas in the ratio of 559 mL of H2 (measured at
STP) for 1.00 g of X reacted. X also combines with
chlorine to form a compound Z, which contains 63.89
percent by mass of chlorine. Deduce the identity of
X.
8.102 A student is given samples of three elements, X, Y,
and Z, which could be an alkali metal, a member of
Group 4A, and a member of Group 5A. She makes
the following observations: Element X has a metallic luster and conducts electricity. It reacts slowly
with hydrochloric acid to produce hydrogen gas.
Element Y is a light-yellow solid that does not conduct electricity. Element Z has a metallic luster and
conducts electricity. When exposed to air, it slowly
forms a white powder. A solution of the white powder in water is basic. What can you conclude about
the elements from these observations?
8.103 Using the following boiling-point data and the procedure in the Chemistry in Action essay on p. 302,
estimate the boiling point of francium:
metal
Li
Na
K Rb
Cs
boiling point (C) 1347 882.9 774 688 678.4
8.104 What is the electron affinity of the Na ion?
8.105 The ionization energies of sodium (in kJ/mol), starting with the first and ending with the eleventh, are
495.9, 4560, 6900, 9540, 13400, 16600, 20120,
25490, 28930, 141360, 170000. Plot the log of ionization energy (y-axis) versus the number of ionization (x-axis); for example, log 495.9 is plotted versus
1 (labeled I1, the first ionization energy), log 4560 is
plotted versus 2 (labeled I2, the second ionization energy), and so on. (a) Label I1 through I11 with the
electrons in orbitals such as 1s, 2s, 2p, and 3s.
(b) What can you deduce about electron shells from
the breaks in the curve?
8.106 Experimentally, the electron affinity of an element
can be determined by using a laser light to ionize the
anion of the element in the gas phase:
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X (g)
h 88n X(g)
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PERIODIC RELATIONSHIPS AMONG THE ELEMENTS
8.107
8.108
8.109
8.110
8.111
8.112
Referring to Table 8.4, calculate the photon wavelength (in nanometers) corresponding to the electron
affinity for chlorine. In what region of the electromagnetic spectrum does this wavelength fall?
Explain, in terms of their electron configurations,
why Fe2 is more easily oxidized to Fe3 than Mn2
to Mn3 .
The standard enthalpy of atomization of an element
is the energy required to convert one mole of an element in its most stable form at 25C to one mole of
monatomic gas. Given that the standard enthalpy of
atomization for sodium is 108.4 kJ/mol, calculate the
energy in kilojoules required to convert one mole of
sodium metal at 25C to one mole of gaseous Na
ions.
Write the formulas and names of the hydrides of the
following second-period elements: Li, C, N, O, F.
Predict their reactions with water.
Based on knowledge of the electronic configuration
of titanium, state which of the following compounds
of titanium is unlikely to exist: K3TiF6, K2Ti2O5,
TiCl3, K2TiO4, K2TiF6.
Name an element in Group 1A or Group 2A that is
an important constituent of each of the following substances: (a) remedy for acid indigestion, (b) coolant
in nuclear reactors, (c) Epsom salt, (d) baking powder, (e) gunpowder, (f) a light alloy, (g) fertilizer that
also neutralizes acid rain, (h) cement, and (i) grit for
icy roads. You may need to ask your instructor about
some of the items.
In halogen displacement reactions a halogen element
can be generated by oxidizing its anions with a halogen element that lies above it in the periodic table.
This means that there is no way to prepare elemental fluorine, since it is the first member of Group 7A.
Indeed, for years the only way to prepare elemental
fluorine was to oxidize F ions by electrolytic means.
Then, in 1986, a chemist reported that by reacting
potassium hexafluoromanganate(IV) (K2MnF6) with
antimony pentafluoride (SbF5) at 150C, he had generated elemental fluorine. Balance the following
equation representing the reaction:
K2MnF6
SbF5 88n KSbF6
MnF3
F2
8.113 Write a balanced equation for the preparation of
(a) molecular oxygen, (b) ammonia, (c) carbon dioxide, (d) molecular hydrogen, (e) calcium oxide.
Indicate the physical state of the reactants and products in each equation.
8.114 Write chemical formulas for oxides of nitrogen with
the following oxidation numbers: 1, 2, 3, 4,
5. (Hint: There are two oxides of nitrogen with 4
oxidation number.)
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8.115 Most transition metal ions are colored. For example,
a solution of CuSO4 is blue. How would you show
that the blue color is due to the hydrated Cu2 ions
2
and not the SO4 ions?
8.116 In general, atomic radius and ionization energy have
opposite periodic trends. Why?
8.117 Explain why the electron affinity of nitrogen is approximately zero, while the elements on either side,
carbon and oxygen, have substantial positive electron
affinities.
8.118 Consider the halogens chlorine, bromine, and iodine.
The melting point and boiling point of chlorine are
101.0C and 34.6C while those of iodine are
113.5C and 184.4C, respectively. Thus chlorine is
a gas and iodine is a solid under room conditions.
Estimate the melting point and boiling point of
bromine. Compare your values with those from a
handbook of chemistry.
8.119 While it is possible to determine the second, third,
and higher ionization energies of an element, the
same cannot usually be done with the electron affinities of an element. Explain.
8.120 The only confirmed compound of radon is radon fluoride, RnF. One reason that it is difficult to study the
chemistry of radon is that all isotopes of radon are
radioactive so it is dangerous to handle the substance.
Can you suggest another reason why there are so few
known radon compounds? (Hint: Radioactive decays
are exothermic processes.)
8.121 Little is known of the chemistry of astatine, the last
member of Group 7A. Describe the physical characteristics that you would expect this halogen to have.
Predict the products of the reaction between sodium
astatide (NaAt) and sulfuric acid. (Hint: Sulfuric acid
is an oxidizing agent.)
8.122 As discussed in the chapter, the atomic mass of argon is greater than that of potassium. This observation created a problem in the early development of
the periodic table because it meant that argon should
be placed after potassium. (a) How was this difficulty
resolved? (b) From the following data, calculate the
average atomic masses of argon and potassium: Ar36 (35.9675 amu; 0.337%), Ar-38 (37.9627 amu;
0.063%), Ar-40 (39.9624 amu; 99.60%); K-39
(38.9637 amu; 93.258%), K-40 (39.9640 amu;
0.0117%), K-41 (40.9618 amu; 6.730%).
8.123 Calculate the maximum wavelength of light (in
nanometers) required to ionize a single sodium atom.
8.124 Predict the atomic number and ground-state electron
configuration of the next member of the alkali metals after francium.
8.125 Why do elements that have high ionization energies
also have more positive electron affinities? Which
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QUESTIONS AND PROBLEMS
group of elements would be an exception to this generalization?
8.126 The first four ionization energies of an element are
approximately 738 kJ/mol, 1450 kJ/mol, 7.7
103 kJ/mol, and 1.1 104 kJ/mol. To which periodic
group does this element belong? Why?
8.127 Some chemists think that helium should properly be
called helon. Why? What does the ending in helium
(-ium) suggest?
8.128 (a) The formula of the simplest hydrocarbon is CH4
(methane). Predict the formulas of the simplest compounds formed between hydrogen and the following
elements: silicon, germanium, tin, and lead.
(b) Sodium hydride (NaH) is an ionic compound.
Would you expect rubidium hydride (RbH) to be
more or less ionic than NaH? (c) Predict the reaction
between radium (Ra) and water. (d) When exposed
to air, aluminum forms a tenacious oxide (Al2O3)
coating that protects the metal from corrosion. Which
metal in Group 2A would you expect to exhibit similar properties? Why?
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8.129 Match each of the elements on the right with its description on the left:
(a) A pale yellow gas
Nitrogen (N2)
that reacts with water.
Boron (B)
(b) A soft metal that reacts
Aluminum (Al)
with water to produce
Fluorine (F2)
hydrogen.
Sodium (Na)
(c) A metalloid that is hard
and has a high melting
point.
(d) A colorless, odorless
gas.
(e) A metal that is more reactive than
iron, but does not corrode
in air.
Answers to Practice Exercises: 8.1 (a) 1s22s22p63s23p64s2,
(b) it is a representative element, (c) diamagnetic. 8.2 Li Be
C. 8.3 (a) Li , (b) Au3 , (c) N3 . 8.4 (a) N, (b) Mg. 8.5 No.
8.6 (a) amphoteric, (b) acidic, (c) basic.
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