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Chapter 7
Periodic Properties of the Elements
Periodic Table
Modern Periodic
Table
– Dmitri Mendeleev
– Elements are
arranged in order of
atomic number
– Organize the table
by chemical
properties
2
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Periodic Table
• Periodicity
– Recurring trends in
the element
properties
• Effective nuclear
charge
– A fundamental
property that leads
to many of the
trends
3
Effective Nuclear Charge
• Properties depend on
attractions between
valence electrons and
the nucleus
• Electrons are both
attracted to the nucleus
and repelled by other
electrons
4
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Effective Nuclear Charge
• Let’s look at an electron’s
attraction to the nucleus
• Effective nuclear charge, Zeff,
– Nuclear charge experienced by
the outermost electrons of an
atom
5
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Effective Nuclear Charge
• Effective nuclear charge, Zeff,
– Nuclear charge experienced
by the outermost electrons
of an atom
– Zeff = Z − S
•Z is the atomic number
•S is a screening constant
–usually close to the number
of inner electrons
•
Effective nuclear charge is a
periodic property:
–
–
Increases across a period
Decreases down a group
6
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Effective Nuclear Charge
• Effective nuclear charge, Zeff,
– Nuclear charge experienced
by the outermost electrons
of an atom
– Zeff = Z − S
•Z is the atomic number
•S is a screening constant
–usually close to the number
of inner electrons
•
Effective nuclear charge is a
periodic property:
–
–
Increases across a period
Decreases down a group
7
http://chem.libretexts.org/@api/deki/files/8872/=Untitled9.jpg?revision=1
Effective Nuclear Charge
• Effective nuclear charge, Zeff,
– Nuclear charge experienced
by the outermost electrons
of an atom
– Zeff = Z − S
•Z is the atomic number
•S is a screening constant
–usually close to the number
of inner electrons
•
Effective nuclear charge is a
periodic property:
–
–
Increases across a period
Decreases down a group
8
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Effective Nuclear Charge
• Effective nuclear charge, Zeff,
– Nuclear charge experienced
by the outermost electrons
of an atom
– Zeff = Z − S
•Z is the atomic number
•S is a screening constant
–usually close to the number
of inner electrons
•
Effective nuclear charge is a
periodic property:
–
–
Increases across a period
Decreases down a group
9
https://classconnection.s3.amazonaws.com/933/flashcards/651080/png/z-eff.png
SIZES OF ATOMS AND IONS
10
Sizes of Atoms
• Atomic radii
– Half of two atoms bond length
— Decrease from left to right across
a period (Zeff ↑)
— Increase from top to bottom of a
group (n ↑)
11
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EXAMPLE ATOMIC SIZE
SKILLBUILDER | Atomic Size
Choose the larger atom in each pair.
(a) Pb or Po
(b) Rb or Na
(c) Sn or Bi
(d) F or Se
Answers:
(a) Pb
(b) Rb
(c)
(d)
12
Ion Size
• Determined by
interatomic distances
in ionic compounds
13
Ion Size
• Determined by
interatomic distances
in ionic compounds
• Ionic size depends on
– the nuclear charge
– the number of
electrons
– the orbitals in which
electrons reside
14
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Ionic Bond: Ion Size
• Cations are smaller
than their parent
atoms:
– Repulsions between
electrons are reduced
• Anions are larger than
their parent atoms:
– Repulsions between
electrons are increased
15
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Ion Size
• Isoelectronic series
– ions have the same number and configuration of electrons
– Ionic size decreases with an increasing nuclear charge
Na+: 1s22s22p6
Ne: 1s22s22p6
Cl-: 1s22s22p63s23p6
Ar: 1s22s22p63s23p6
16
Ion Size
• Isoelectronic series
– ions have the same number and configuration of electrons
– Ionic size decreases with an increasing nuclear charge
 An Isoelectronic Series (10 electrons)
O2–
1.26 Å
F–
1.19 Å
Na+
1.16 Å
Mg2+
0.86 Å
Al3+
0.68 Å
• Note increasing nuclear charge with decreasing ionic
radius as atomic number increases
17
EXAMPLE ATOMIC SIZE
Choose the larger atom in each pair.
(a) C or O
(b) Li or K
(c) C or Al
(d) Se or I
SOLUTION
(a)
C or O
Carbon atoms are larger than O
atoms because, as you trace the
path between C and O on the
periodic table, you move to the
right within the same period.
Atomic size decreases as you go to
the right.
18
EXAMPLE ATOMIC SIZE
Choose the larger atom in each pair.
(a) C or O
(b) Li or K
(c) C or Al
(d) Se or I
SOLUTION
(b)
Li or K
Potassium atoms are larger than
Li atoms because, as you trace
the path between Li and K on the
periodic table, you move down a
column. Atomic size increases as
you go down a column.
19
EXAMPLE ATOMIC SIZE
Choose the larger atom in each pair.
(a) C or O
(b) Li or K
(c) C or Al
(d) Se or I
20
EXAMPLE ATOMIC SIZE
Choose the larger atom in each pair.
(a) C or O
(b) Li or K
(c) C or Al
(d) Se or I
21
IONIZATION ENERGY
22
Ionization Energy (I)
• Ionization energy
– Minimum energy required to
remove an electron from the
ground state of a gaseous
atom or ion
– First ionization energy is that
energy required to remove
the first electron
– Second ionization energy is
that energy required to
remove the second electron,
etc
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https://www.thunderbolts.info/wp/wp-content/uploads/2011/08/ionization-image.jpg
Ionization Energy
• Requires more energy to remove each successive electron
• Higher the ionization energy, the more difficult it is to remove an
electron
• When all valence electrons have been removed, it takes a
great deal more energy to remove the next electron
24
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Periodic Trends in
First Ionization Energy (I1)
•
•
•
•
I1 generally
increases
across a
period
I1 generally
increases up
a group
The s- and
p-block
elements show a larger
elements
show aof values for I
range
1
The d-block generally increases slowly across the period
25
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Factors that Influence Ionization
Energy
•
•
Smaller atoms
have higher I
values
I values
depend on
effective
nuclear charge
and average
distance of the
electron from
the nucleus
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First Ionization Energy: Irregularities
• Trend is not followed when the added valence
electron
– enters a new sublevel (higher energy subshell)
– is the first electron to pair in one orbital of the sublevel
(electron repulsions lower energy)
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Example First Ionization Energy
On the basis of periodic trends, determine the element with the higher first
ionization energy in each pair (if possible).
(a) Al or S
(b) As or Sb
(c) N or Si
(d) O or Cl
Solution
(a)
Al or S
S has a higher first ionization
energy than Al because, as you trace
the path between Al and S on the
periodic table, you move to the right
within the same row. Ionization
energy increases as you go to the
right because of increasing effective
nuclear charge.
Example First Ionization Energy
On the basis of periodic trends, determine the element with the higher first
ionization energy in each pair (if possible).
(a) Al or S
(b) As or Sb
(c) N or Si
(d) O or Cl
Solution
(b)
As or Sb
As has a higher first ionization
energy than Sb because, as you
trace the path between As and Sb
on the periodic table, you move
down a column. Ionization energy
decreases as you go down a
column because of the increasing
size of orbitals with increasing n.
Example First Ionization Energy
On the basis of periodic trends, determine the element with the higher first
ionization energy in each pair (if possible).
(a) Al or S
(b) As or Sb
(c) N or Si
(d) O or Cl
Example First Ionization Energy
On the basis of periodic trends, determine the element with the higher first
ionization energy in each pair (if possible).
(a) Al or S
(b) As or Sb
(c) N or Si
(d) O or Cl
ELECTRON AFFINITY
32
Electron Affinity
• Electron affinity
– Energy change accompanying the addition of
an electron to a gaseous atom
Cl + e−  Cl−
– Typically exothermic, so negative value
33
General Trend in Electron Affinity
• Trends
– Not much change in a
group
– Increases across a
period
– Three notable
exceptions include:
1)
2)
3)
Group 2A: s sublevel
is full!
Group 5A: p sublevel
is half-full!
Group 8A: p sublevel
is full!
34
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The correct order of
increasing electron affinity
(most negative  least
negative) is
a.
b.
c.
d.
O < Cl < B < C.
O < Cl < C < B.
Cl < O < C < B.
Cl < O < B < C.
The correct order of
increasing electron affinity
(most negative  least
negative) is
a.
b.
c.
d.
O < Cl < B < C.
O < Cl < C < B.
Cl < O < C < B.
Cl < O < B < C.
METALS, NONMETALS, AND
METALLOIDS
37
Metal, Nonmetals, and Metalloids
38
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Metal, Nonmetals, and Metalloids
Properties of metals:
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except
mercury)
Low ionization energies/form cations easily
39
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Metal, Nonmetals, and Metalloids
Properties of nonmetals:
Solid, liquid, or gas (depends on element)
Solids are dull, brittle, poor conductors
Large negative electronegativity
40
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Metal, Nonmetals, and Metalloids
Properties of metalloids:
Characteristics of metals and nonmetals
Several metalloids are electrical
semiconductors (computer chips)
41
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Metalloids
• Metalloids have some characteristics of metals and
some of nonmetals.
• Several metalloids are electrical semiconductors
(computer chips).
42
Metals vs Nonmetals
• Metals tend to form cations
• Nonmetals tend to form anions
43
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Metal Chemistry
• Compounds the contain metals and nonmetals
tend to be ionic
• Metal oxides tend to be basic
44
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Nonmetal Chemistry
• Substances containing only nonmetals are
molecular compounds
• Most nonmetal oxides are acidic
45
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EXAMPLE METALLIC CHARACTER
Choose the more metallic element from each pair.
(a) Sn or Te
(b) Si or Sn
(c) Br or Te
(d) Se or I
SOLUTION
(a)
Sn or Te
Sn is more metallic than Te
because, as you trace the path
between Sn and Te on the periodic
table, you move to the right within
the same period. Metallic character
decreases as you go to the right.
46
EXAMPLE METALLIC CHARACTER
Choose the more metallic element from each pair.
(a) Sn or Te
(b) Si or Sn
(c) Br or Te
(d) Se or I
SOLUTION
(b)
Si or Sn
Sn is more metallic than Si
because, as you trace the path
between Si and Sn on the
periodic table, you move down a
column. Metallic character
increases as you go down a
column.
47
EXAMPLE METALLIC CHARACTER
Choose the more metallic element from each pair.
(a) Sn or Te
(b) Si or Sn
(c) Br or Te
(d) Se or I
48
EXAMPLE METALLIC CHARACTER
Choose the more metallic element from each pair.
(a) Sn or Te
(b) Si or Sn
(c) Br or Te
(d) Se or I
49
Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
50
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Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
– Only found in compounds
in nature
51
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Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
– Only found in compounds
in nature
– Low melting points
52
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Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
– Only found in compounds
in nature
– Low melting points
– Low densities
53
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Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
– Only found in compounds
in nature
– Low melting points
– Low densities
– Low ionization energies
54
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Alkali Metals
• Alkali metals
– Soft
– Metallic solids
• Typical metallic properties
(luster, conductivity)
– Only found in compounds
in nature
– Low densities
– Low melting points
– Low ionization energies
– Reactions with water are
exothermic
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55
Alkali Metals
• Alkali metals
– Emit characteristic color in
a flame
• Flame test: Qualitative test
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56
Alkali Metal Chemistry
• Lithium reacts with oxygen to make an oxide:
4 Li + O2  2 Li2O
• Sodium reacts with oxygen to form a peroxide:
2 Na + O2  Na2O2
• K, Rb, and Cs also form superoxides:
M + O2  MO2
57
Alkaline Earth Metals
• Alkaline Earth Metals
– Reaction with water
• Beryllium does not react
with water
• Magnesium reacts
readily with steam
• Other alkaline earth
metals react readily with
water
– Reactivity tends to
increase down the group
58
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Alkaline Earth Metals
• Alkaline Earth Metals
– Reaction with water
• Beryllium does not react
with water
• Magnesium reacts
readily with steam
• Other alkaline earth
metals react readily with
water
– Reactivity tends to
increase down the group
– Emit characteristic color
in a flame
https://qph.is.quoracdn.net/main-qimg-3337403d76081268fb74f7351c1c7ec5?convert_to_webp=true
59
Alkaline Earth Metals
• Alkaline Earth Metals
– Higher densities and melting points than alkali metals
– Low ionization energies
• not as low as alkali metals
60
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Chalcogens
• Oxygen, sulfur, and selenium are nonmetals
• Tellurium is a metalloid
• The radioactive polonium is a metal
61
Halogens
• Halogens are typical nonmetals
• Have highly negative electron affinities
– Exist as anions in nature
• React with metals to form metal halides
62
Noble Gases
• Noble gases have
–
–
–
–
Very large ionization energies
Positive electron affinities (can’t form stable anions)
Relatively unreactive
Found as monatomic gases
63
64
EXAMPLE METALLIC CHARACTER
Continued
SKILLBUILDER | Metallic Character
Choose the more metallic element from each pair.
(a) Ge or In
(b) Ga or Sn
(c) P or Bi
(d) Bor N
Answers:
65
The first Periodic Table was
created by
a.
b.
c.
d.
Amadeo Avogadro.
Dmitri Mendeleev.
Henry Moseley.
Ernest Rutherford.
66
The effective nuclear charge
felt by an atom’s valence
electrons (X) going from left to
right and (Y) going down a
column on the periodic table.
a.
b.
c.
d.
X = increases
X = increases
X = decreases
X = decreases
Y = increases
Y = decreases
Y = increases
Y = decreases
67
True or False:
The large number of valence
electrons in a chlorine atom
greatly reduces the effective
nuclear charge of the atom.
a.True
b.False
68
The atom with the smallest
atomic radius is
a.
b.
c.
d.
Ca.
Cr.
Co.
Br.
69
The correct order of
increasing atomic radius
(smallest  largest) is
a.
b.
c.
d.
Na < Mg < K < Rb.
Mg < Na < K < Rb.
Rb < K < Na < Mg.
Rb < K < Mg < Na.
70
The shortest distance between
the nuclei of two identical atoms
when they collide is called the
___ radius.
a.
b.
c.
d.
anionic
cationic
effective
van der Waals
71
The statements below refer to
ionic radii. Which statement is
FALSE?
a.
b.
c.
d.
Br1– is larger than Cl1–.
Se2– is larger than Br1–.
K1+ is larger than Ca2+.
Na1+ is larger than K1+.
72
Two ions are isoelectronic if
they have the same
a.
b.
c.
d.
charge.
number of protons.
number of electrons.
number of neutrons.
73
Which set represents an
isoelectronic series?
a.
b.
c.
d.
Ne, Ar, Kr, Xe
Al, Si; Ge, As; Sb, Te
Se2−, Br−, Rb+, Sr2+
Cu 2+, Ag+, Au
74
When a main-group metal ion
loses an electron to form a
cation, the ionic radius is ___
the atomic radius.
a.
b.
c.
d.
greater than
less than
the same as
impossible to predict in relation to
75
The minimum energy needed
to remove an electron from an
atom in its ground state is
called the atom’s
a.
b.
c.
d.
potential energy.
activation energy.
electron affinity.
ionization energy.
76
The energy change that
occurs when an electron is
added to an atom is called
the atom’s
a.
b.
c.
d.
potential energy.
activation energy.
electron affinity.
ionization energy.
77
The correct order of
increasing first ionization
energy (smallest  largest) is
a.
b.
c.
d.
Na < Ca < Al < Sn.
Na < Al < Ca < Sn.
Na < Al < Sn < Ca.
Ca < Na < Sn < Al.
78
The atom with successive
ionization energies of
578, 1817, 2745, and
11,577 kJ/mol is
a.
b.
c.
d.
Al.
Na.
Mg.
P.
79
The correct order of
increasing electron affinity
(most negative  least
negative) is
a.
b.
c.
d.
O < Cl < B < C.
O < Cl < C < B.
Cl < O < C < B.
Cl < O < B < C.
80
Most metal oxides form
_______ solutions when
dissolved in water.
a.
b.
c.
d.
acidic
basic
neutral
amphoteric
81
Most nonmetal oxides form
_______ solutions when
dissolved in water.
a.
b.
c.
d.
acidic
basic
neutral
amphoteric
82
The metalloid with the least
metallic character is
a.
b.
c.
d.
e.
As.
B.
Ge.
Sb.
Si.
83
A soft gray solid reacts with
water to form a flammable gas
and a basic solution. The solid
is most likely to be a
a.
b.
c.
d.
noble gas.
halogen.
group IIA metal.
transition metal.
84
Ozone is an allotrope of which
element?
a.
b.
c.
d.
hydrogen
oxygen
sulfur
chlorine
85
Unknown solid X turns a
Bunsen burner flame green
and produces a purple color
on reaction with NaOCl. X is
a.
b.
c.
d.
BaI2.
BaBr2.
SrI2.
SrBr2.
86
The noble gases
a.
b.
c.
d.
are monatomic.
have filled s and p subshells.
are generally unreactive.
All of the above.
87