Chemistry
Third Edition
Julia Burdge
Lecture PowerPoints
Chapter 7
Electron Configuration
and the Periodic Table
Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display.
CHAPTER
7.1
7.2
7.3
7.4
7.5
7.6
7.7
7
Electron Configuration and the
Periodic Table
Development of the Periodic Table
The Modern Periodic Table
Effective Nuclear Charge
Periodic Trends in Properties of Elements
Electron Configuration of Ions
Ionic Radius
Periodic Trends in Chemical Properties of the Main
Group Elements
2
7.1
Development of the Periodic Table
Topics
Development of the Periodic Table
3
7.1
Development of the Periodic Table
Development of the Periodic Table
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 Newlands’s work was not accepted by
the scientific community.
4
7.1
Development of the Periodic Table
Development of the Periodic Table
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—a
phenomenon known as periodicity.
5
7.1
Development of the Periodic Table
Development of the Periodic Table
Mendeleev ordered the elements based on atomic mass,
which led to inconsistencies.
For example, argon and potassium would switch places on the
periodic table if it were ordered by atomic mass.
In such a case, potassium (very reactive) would be grouped
with neon and helium (unreactive) and argon (unreactive)
would be grouped with sodium and lithium (very reactive).
This seriously upsets the idea of periodicity!
6
7.1
Development of the Periodic Table
Development of the Periodic Table
The key to periodicity turns out to be atomic number instead
of atomic mass.
Modern periodic table orders the elements by atomic
number.
Elements in a column of the modern periodic table have very
similar chemical properties.
7
SAMPLE PROBLEM
7.1
What elements would you expect to exhibit properties most
similar to those of chlorine?
Strategy
Because elements in the same group tend to have similar
properties, you should identify elements in the same group as
chlorine.
Setup
Chlorine is a member of Group 7A.
8
SAMPLE PROBLEM
7.1
Solution
Fluorine, bromine, and iodine, the other nonmetals in Group
7A, should have properties most similar to those of chlorine.
9
7.2
The Modern Periodic Table
Topics
Classification of Elements
Representing Free Elements in Chemical Equations
10
7.2
The Modern Periodic Table
Classification of Elements
Based on the type of subshell containing the outermost
electrons, the elements can be divided into categories—the
main group elements, the noble gases, the transition
elements (or transition metals), the lanthanides, and the
actinides.
The main group elements (also called the representative
elements) are the elements in Groups 1A through 7A.
11
7.2
The Modern Periodic Table
12
7.2
The Modern Periodic Table
Classification of Elements
The outermost electrons of an atom are called valence
electrons, which are the ones involved in the formation of
chemical bonds between atoms.
The similarity of the valence electron configurations (i.e., they
have the same number and type of valence electrons) is what
makes the elements in the same group resemble one another
chemically.
For instance, the halogens (Group 7A) all have outer electron
configurations of ns2np5, and they have similar properties.
13
7.2
The Modern Periodic Table
Classification of Elements
14
7.2
The Modern Periodic Table
Representing Free Elements in Chemical Equations
Metals
Because metals typically do not exist in discrete molecular
units but rather in complex, three-dimensional networks of
atoms, 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.
15
7.2
The Modern Periodic Table
Representing Free Elements in Chemical Equations
Nonmetals
For nonmetals that exist as polyatomic molecules, we
generally use the molecular formula in equations: H2, N2, O2,
F2, Cl2, Br2, I2, and P4, for example.
In the case of sulfur, however, we usually use the empirical
formula S rather than the molecular formula S8.
16
7.2
The Modern Periodic Table
Representing Free Elements in Chemical Equations
Noble Gases
All the noble gases exist as isolated atoms, so we use their
symbols: He, Ne, Ar, Kr, Xe, and Rn.
Metalloids
The metalloids, like the metals, all have complex threedimensional networks, so we also represent them with their
empirical formulas—that is, their symbols: B, Si, Ge, and so
on.
17
7.3
Effective Nuclear Charge
Topics
Effective Nuclear Charge
18
7.3
Effective Nuclear Charge
Effective Nuclear Charge
Nuclear charge (Z) is simply the number of protons in the
nucleus of an atom.
An electron in a many-electron atom is partially shielded from
the positive charge of the nucleus by the other electrons in
the atom.
Effective nuclear charge (Zeff) is the actual magnitude of
positive charge that is “experienced” by a shielded electron in
the atom. Because of shielding, the effective nuclear charge is
less than the actual nuclear charge.
19
7.3
Effective Nuclear Charge
Effective Nuclear Charge
Although all the electrons in an atom shield one another to
some extent, those that are most effective at shielding are the
core electrons.
As a result, the value of Zeff increases steadily from left to
right across a period of the periodic table because the
number of core electrons remains the same (only the number
of protons, Z, and the number of valence electrons increases).
20
7.3
Effective Nuclear Charge
Effective Nuclear Charge
In general, the effective nuclear charge is given by
where  is the shielding constant. The shielding constant is
greater than zero but smaller than Z.
21
7.3
Effective Nuclear Charge
Effective Nuclear Charge
The change in Zeff as we move from the top of a group to the
bottom is generally less significant than the change as we
move across a period.
Although each step down a group represents a large increase
in the nuclear charge, there is also an additional shell of core
electrons to shield the valence electrons from the nucleus.
Consequently, the effective nuclear charge changes less than
the nuclear charge as we move down a column of the
periodic table.
22
7.4
Periodic Trends in Properties of Elements
Topics
Atomic Radius
Ionization Energy
Electron Affinity
Metallic Character
Explaining Periodic Trends
23
7.4
Periodic Trends in Properties of Elements
Atomic Radius
There are two ways in which the atomic radius is commonly
defined.
One is the metallic radius, which is half
the distance between the nuclei of two
adjacent, identical metal atoms.
24
7.4
Periodic Trends in Properties of Elements
Atomic Radius
The other is the covalent radius, which is half the distance
between adjacent, identical nuclei in a molecule.
25
7.4
Periodic Trends in Properties of Elements
Atomic Radius
The atomic radius decreases as we move from left to right
across a period and increases from top to bottom as we move
down within a group.
26
7.4
Periodic Trends in Properties of Elements
27
7.4
Periodic Trends in Properties of Elements
Atomic Radius
The increase down a group is fairly easily explained.
As we step down a column, the outermost occupied shell has
an ever-increasing value of n, so it lies farther from the
nucleus, making the radius bigger.
28
7.4
Periodic Trends in Properties of Elements
Atomic Radius
As we move from left to right across a period, the effective
nuclear charge increases and each step to the right adds
another electron to the valence shell.
So, there will be a more powerful attraction between the
nucleus and the valence shell when the magnitudes of both
charges increase.
The result is that as we step across a period the valence shell
is drawn closer to the nucleus, making the atomic radius
smaller.
29
SAMPLE PROBLEM
7.3
Referring only to a periodic table, arrange the elements P, S,
and O in order of increasing atomic radius.
Strategy
Use the left-to-right (decreasing) and top-to-bottom
(increasing) trends to compare the atomic radii of two of the
three elements at a time.
Setup
Sulfur is to the right of phosphorus in the third row, so sulfur
should be smaller than phosphorus. Oxygen is above sulfur in
Group 6A, so oxygen should be smaller than sulfur.
30
SAMPLE PROBLEM
Solution
7.3
O < S < P.
31
7.4
Periodic Trends in Properties of Elements
Ionization Energy
Ionization energy (IE) is the minimum energy required to
remove an electron from an atom in the gas phase.
Sodium, for example, has an ionization energy of 495.8
kJ/mol, which is the energy input required to drive the
process
Specifically, this is the first ionization energy of sodium,
IE1(Na), which corresponds to the removal of the most loosely
held electron.
32
7.4
Periodic Trends in Properties of Elements
Ionization Energy
33
7.4
Periodic Trends in Properties of Elements
Ionization Energy
34
7.4
Periodic Trends in Properties of Elements
Ionization Energy
In general, as effective nuclear charge
increases, ionization energy also
increases.
Thus, IE1 increases from left to right
across a period.
Despite this trend, IE1 for a Group 3A
element is smaller than that for the
corresponding Group 2A element.
35
7.4
Periodic Trends in Properties of Elements
Ionization Energy
Likewise, IE1 for a Group 6A element
is smaller than that for the
corresponding Group 5A element.
Both of these interruptions of the
upward trend in IE1 can be explained
by using electron configuration.
36
7.4
Periodic Trends in Properties of Elements
Ionization Energy
37
7.4
Periodic Trends in Properties of Elements
Ionization Energy
38
SAMPLE PROBLEM
7.4
Would you expect Na or Mg to have the greater first
ionization energy (IE1)? Which should have the greater second
ionization energy (IE2)?
Strategy
Consider effective nuclear charge and electron configuration
to compare the ionization energies.
Effective nuclear charge increases from left to right in a period
(thus increasing IE), and it is more difficult to remove a paired
core electron than an unpaired valence electron.
39
SAMPLE PROBLEM
7.4
Setup
Na is in Group 1A, and Mg is beside it in Group 2A. Na has
one valence electron, and Mg has two valence electrons.
Solution
IE1(Mg) > IE1(Na) because Mg is to the right of Na in the
periodic table (i.e., Mg has the greater effective nuclear
charge, so it is more difficult to remove its electron).
IE2(Na) > IE2(Mg) because the second ionization of Mg
removes a valence electron, whereas the second ionization of
Na removes a core electron.
40
7.4
Periodic Trends in Properties of Elements
Electron Affinity
Electron affinity (EA) is the energy released (the negative of
the enthalpy change H) when an atom in the gas phase
accepts an electron.
41
7.4
Periodic Trends in Properties of Elements
Electron Affinity
42
7.4
Periodic Trends in Properties of Elements
43
7.4
Periodic Trends in Properties of Elements
Electron Affinity
Like ionization energy, electron affinity increases from left to
right across a period.
This trend in EA is due to the increase in effective nuclear
charge from left to right (i.e., it becomes progressively easier
to add a negatively charged electron as the positive charge of
the element’s nucleus increases).
44
7.4
Periodic Trends in Properties of Elements
Electron Affinity
There are also periodic interruptions of the upward trend of
EA from left to right, similar to those observed for IE1,
although they do not occur for the same elements.
45
7.4
Periodic Trends in Properties of Elements
Electron Affinity
46
SAMPLE PROBLEM
7.5
For each pair of elements, indicate which one you would
expect to have the greater first electron affinity, EA1:
(a) Al or Si, (b) Si or P.
47
SAMPLE PROBLEM
7.5
Strategy
Consider the effective nuclear charge and electron
configuration to compare the electron affinities.
The effective nuclear charge increases from left to right in a
period (thus generally increasing EA), and it is more difficult
to add an electron to a partially occupied orbital than to an
empty one.
Writing out orbital diagrams for the valence electrons is
helpful for this type of problem.
48
SAMPLE PROBLEM
7.5
Setup
49
SAMPLE PROBLEM
7.5
Solution
(a) EA1(Si)> EA1(Al) because Si is to the right of Al and
therefore has a greater effective nuclear charge.
(b) EA1(Si) > EA1(P) because although P is to the right of Si in
the third period of the periodic table (giving P the larger
Zeff), adding an electron to a P atom requires placing it in a
3p orbital that is partially occupied.
The energy cost of pairing electrons outweighs the energy
advantage of adding an electron to an atom with a larger
effective nuclear charge.
50
7.4
Periodic Trends in Properties of Elements
Metallic Character
Metals tend to
• Be shiny, lustrous, and malleable
• Be good conductors of both heat and electricity
• Have low ionization energies (so they commonly form
cations)
• Form ionic compounds with chlorine (metal chlorides)
• Form basic, ionic compounds with oxygen (metal oxides)
51
7.4
Periodic Trends in Properties of Elements
Metallic Character
Nonmetals tend to
• Vary in color and lack the shiny appearance associated
with metals
• Be brittle, rather than malleable
• Be poor conductors of both heat and electricity
• Form acidic, molecular compounds with oxygen
• Have high electron affinities (so they commonly form
anions)
52
7.4
Periodic Trends in Properties of Elements
Metallic Character
Metallic character increases from top to bottom in a group
and decreases from left to right within a period.
Metalloids are elements with properties intermediate
between those of metals and nonmetals.
53
7.5
Electron Configuration of Ions
Topics
Ions of Main Group Elements
Ions of d-Block Elements
54
7.5
Electron Configuration of Ions
Ions of Main Group Elements
The 1s2 configuration of He and the ns2np6 (n  2) valence
electron configurations of the other noble gases are
extraordinarily stable.
Other main group elements tend to either lose or gain the
number of electrons needed to achieve the same number of
electrons as the nearest noble gas.
Species with identical electron configurations are called
isoelectronic.
55
7.5
Electron Configuration of Ions
Ions of Main Group Elements
To write the electron configuration of an ion formed by a
main group element, we first write the configuration for the
atom and either add or remove the appropriate number of
electrons.
Electron configurations for the sodium and chloride ions are
56
7.5
Electron Configuration of Ions
Ions of Main Group Elements
We can also write electron configurations for ions using the
noble gas core.
57
SAMPLE PROBLEM
7.7
Write electron configurations for the following ions of main
group elements:
(a) N3–, (b) Ba2+, and (c) Be2+.
Strategy
First write electron configurations for the atoms. Then add
electrons (for anions) or remove electrons (for cations) to
account for the charge.
58
SAMPLE PROBLEM
7.7
Setup
(a) N3– forms when N ([He]2s22p3), a main group nonmetal,
gains three electrons.
(b) Ba2+ forms when Ba ([Xe]6s2) loses two electrons.
(c) Be2+ forms when Be ([He]2s2) loses two electrons.
Solution
(a) [Ne]
(b) [Xe]
(c) [He]
59
7.5
Electron Configuration of Ions
Ions of d-Block Elements
We might expect the two electrons lost in the formation of
the Fe2+ ion to come from the 3d subshell.
It turns out, though, that an atom always loses electrons first
from the shell with the highest value of n.
In the case of Fe, that would be the 4s subshell.
60
7.6
Ionic Radius
Topics
Comparing Ionic Radius with Atomic Radius
Isoelectronic Series
61
7.6
Ionic Radius
Comparing Ionic Radius with Atomic Radius
When an atom gains or loses one or more electrons to
become an ion, its radius changes.
The ionic radius, the radius of a cation or an anion, affects the
physical and chemical properties of an ionic compound.
The three-dimensional structure of an ionic compound, for
example, depends on the relative sizes of its cations and
anions.
62
7.6
Ionic Radius
Comparing Ionic Radius with Atomic Radius
When an atom loses an electron and becomes a cation, its
radius decreases due in part to a reduction in electronelectron repulsions (and consequently a reduction in
shielding) in the valence shell.
A significant decrease in radius occurs when all of an atom’s
valence electrons are removed.
This is the case with ions of most main group elements, which
are isoelectronic with the noble gases preceding them.
63
7.6
Ionic Radius
Comparing Ionic Radius with Atomic Radius
64
7.6
Ionic Radius
Isoelectronic Series
An isoelectronic series is a series of two or more species that
have identical electron configurations, but different nuclear
charges.
For example, O2–, F–, and Ne constitute an isoelectronic
series.
Although these three species have identical electron
configurations, they have different radii.
65
7.6
Ionic Radius
Isoelectronic Series
In an isoelectronic series, the species with the smallest
nuclear charge (i.e., the smallest atomic number Z) will have
the largest radius.
The species with the largest nuclear charge (i.e., the largest Z)
will have the smallest radius.
66
7.6
Ionic Radius
67
SAMPLE PROBLEM
7.9
Identify the isoelectronic series in the following group of
species, and arrange the ions in order of increasing radius:
K+, Ne, Ar, Kr, P3–, S2–, and Cl–.
68
SAMPLE PROBLEM
7.9
Setup
The number of electrons in each species is as follows:
18 (K+), 10 (Ne), 18 (Ar), 36 (Kr), 18 (P3–), 18 (S2–), and 18 (Cl–).
The species with 18 electrons constitute the isoelectronic
series.
The nuclear charges of the ions with 18 electrons are +19 (K+),
+15 (P3–), +16 (S2–), and +17 (Cl–).
69
SAMPLE PROBLEM
7.9
Solution
The isoelectronic series includes K+, Ar, P3–, S2–, and Cl–.
The ions, in order of increasing radius, are:
K+ < Cl– < S2– < P3–
70
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Topics
General Trends in Chemical Properties
Properties of the Active Metals
Properties of Other Main Group Elements
Comparison of Group 1A and Group 1B Elements
Variation in Properties of Oxides Within a Period
71
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
General Trends in Chemical Properties
We have said that elements in the same group resemble one
another in chemical behavior because they have similar
valence electron configurations.
This statement, although correct in the general sense, must
be applied with caution.
Chemists have long known that the properties of the first
member of each group (Li, Be, B, C, N, O, and F) are different
from those of the rest of the members of the same group.
72
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
General Trends in Chemical Properties
Another trend in the chemical behavior of main group
elements is the diagonal relationship.
Diagonal relationships refer to similarities between pairs of
elements in different groups and periods of the periodic
table.
73
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
General Trends in Chemical Properties
The reason for this phenomenon is the similarity of charge
densities of their cations.
74
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
General Trends in Chemical Properties
Hydrogen (1s1)
There is no completely suitable position for hydrogen in the
periodic table (it really belongs in a group by itself).
75
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 1A Elements (ns1, n ≥ 2)
These elements all have low ionization energies, making it
easy for them to become M+ ions.
In fact, these metals are so reactive that they are never found
in nature in the pure elemental state.
They react with water to produce hydrogen gas and the
corresponding metal hydroxide:
76
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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77
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 1A Elements (ns1, n ≥ 2)
Lithium forms lithium oxide (containing the oxide ion, O2–):
The other alkali metals all form oxides or peroxides
containing the peroxide ion, O22–):
78
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 1A Elements (ns1, n ≥ 2)
Potassium, rubidium, and cesium also form superoxides
(containing the superoxide ion, O2– ):
79
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 2A Elements (ns2, n ≥ 2)
As a group, the alkaline earth metals are somewhat less
reactive than the alkali metals.
Both the first and the second ionization energies decrease
(and metallic character increases) from beryllium to barium.
Group 2A elements tend to form M2+ ions, where M denotes
an alkaline earth metal atom.
80
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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81
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 2A Elements (ns2, n ≥ 2)
The reactions of alkaline earth metals with water vary
considerably.
Beryllium does not react with water; magnesium reacts slowly
with steam; and calcium, strontium, and barium react
vigorously with cold water.
82
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of the Active Metals
Group 2A Elements (ns2, n ≥ 2)
The reactivity of the alkaline earth metals toward oxygen also
increases 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.
83
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 3A Elements (ns2np1, n ≥ 2)
Boron, the first member of the group, is a metalloid; the
others (Al, Ga, In, and Tl) are metals.
Boron does not form binary ionic compounds and is
unreactive toward both oxygen and water.
Aluminum, the next element in the group, readily forms
aluminum oxide when exposed to air:
84
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 3A Elements (ns2np1, n ≥ 2)
Aluminum forms the Al3+ ion. It reacts with hydrochloric acid
according to the equation:
The other Group 3A metals (Ga, In, and Tl) can form both M+
and M3+ ions.
As we move down the group, the M+ ion becomes the more
stable of the two.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 4A Elements (ns2np2, n ≥ 2)
Carbon, the first member of the group, is a nonmetal,
whereas silicon and germanium, the next two members, are
metalloids.
Tin and lead, the last two members of the group, are metals.
They do not react with water, but they do react with aqueous
acid to produce hydrogen gas:
87
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 4A Elements (ns2np2, n ≥ 2)
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.
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 the more stable
one.
88
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 4A Elements (ns2np2, n ≥ 2)
The outer electron configuration of lead is 6s26p2, and lead
tends to lose only the 6p electrons to form Pb2+ rather than
both the 6p and 6s electrons to form Pb4+.
89
7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 5A Elements (ns2np3, n ≥ 2)
Nitrogen and phosphorus are nonmetals, arsenic and
antimony are metalloids, and bismuth is a metal.
Because Group 5A contains elements in all three categories,
we expect greater variation in their chemical properties.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 5A Elements (ns2np3, n ≥ 2)
Elemental nitrogen is a diatomic gas (N2).
It forms a variety of oxides (NO, N2O, NO2, N2O4, and N2O5), all
of which are gases except for N2O5, which is a solid at room
temperature.
Nitrogen has a tendency to accept three electrons to form the
nitride ion (N3–). Most metal nitrides, such as Li3N and Mg3N2,
are ionic compounds.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 5A Elements (ns2np3, n ≥ 2)
Phosphorus exists as individual P4 molecules (white
phosphorus) or chains of P4 molecules (red phosphorus).
It forms two solid oxides with the formulas P4O6 and P4O10.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 6A Elements (ns2np4, n ≥ 2)
The first three members of the group (oxygen, sulfur, and
selenium) are nonmetals, whereas the last two (tellurium and
polonium) are metalloids.
Oxygen is a colorless, odorless, diatomic gas; elemental sulfur
and selenium exist as the molecules S8 and Se8, respectively;
and tellurium and polonium have more extensive threedimensional structures.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 6A Elements (ns2np4, n ≥ 2)
Oxygen has a tendency to accept two electrons to form the
oxide ion (O2–) in many compounds.
Sulfur, selenium, and tellurium also form ions by accepting
two electrons: S2–, Se2–, and Te2–.
The elements in Group 6A (especially oxygen) form a large
number of molecular compounds with nonmetals.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 7A Elements (ns2np5, n ≥ 2)
All the halogens are nonmetals with the general formula X2,
where X denotes a halogen element.
Like the Group 1A metals, the Group 7A nonmetals are too
reactive to be found in nature in the elemental form.
The halogens have high ionization energies and large,
energetically favorable electron affinities.
Anions derived from the halogens (F–, Cl–, Br–, and I–) are
called halides.
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7.7
Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 7A Elements (ns2np5, n ≥ 2)
The vast majority of alkali metal halides are ionic
compounds.
The halogens also form many molecular compounds among
themselves, such as ICl and BrF3, and with nonmetals in
other groups, such as NF3, PCl5, and SF6.
The halogens react with hydrogen to form hydrogen halides:
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Periodic Trends in Chemical Properties of the
Main Group Elements
© The McGraw-Hill Companies,
Inc./Stephen Frisch, photographer
The McGraw-Hill Companies, Inc./Stephen
Frisch, photographer
© Charles D. Winters/Photo Researchers
© The McGraw-Hill Companies,
Inc./Charles D. Winters, photographer
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Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 8A Elements (ns2np6, n ≥ 2)
All the noble gases exist as monatomic species.
With the exception of helium, which has the electron
configuration 1s2, their atoms have completely filled outer ns
and np subshells.
Their electron configurations give the noble gases their great
stability.
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Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 8A Elements (ns2np6, n ≥ 2)
The Group 8A ionization energies are among the highest of all
the elements.
Their electron affinities are all less than zero, so they have no
tendency to accept extra electrons.
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Periodic Trends in Chemical Properties of the
Main Group Elements
Properties of Other Main Group Elements
Group 8A Elements (ns2np6, n ≥ 2)
Beginning in 1963, compounds were prepared from the
heavier members of the group by exposing them to very
strong oxidizing agents such as fluorine and oxygen.
Some of the compounds that have been prepared are XeF4,
XeO3, XeOF4, KrF2, and most recently, HArF.
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Periodic Trends in Chemical Properties of the
Main Group Elements
The McGraw-Hill Companies, Inc./Stephen
Frisch, photographer
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Periodic Trends in Chemical Properties of the
Main Group Elements
Comparison of Group 1A and Group 1B Elements
Although the outer electron configurations of Groups 1A and
1B are similar (members of both groups have a single
valence electron in an s orbital), their chemical properties
are very different.
The first ionization energies of Cu, Ag, and Au are 745, 731,
and 890 kJ/mol, respectively.
Because these values are considerably larger than those of
the alkali metals, the Group 1B elements are much less
reactive.
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Periodic Trends in Chemical Properties of the
Main Group Elements
Comparison of Group 1A and Group 1B Elements
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 by
the completely filled noble gas cores).
Consequently, the outer s electrons of the Group 1B
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.
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Periodic Trends in Chemical Properties of the
Main Group Elements
Variation in Properties of Oxides Within a Period
Most oxides can be classified as acidic or basic depending on
whether they produce acidic or basic solutions when
dissolved in water (or whether they react as acids or bases).
Some oxides are amphoteric, which means that they display
both acidic and basic properties.
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