Electronic Structure and the
Periodic Table
CHAPTER 6
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Chemical
Periodicity
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The electron configuration of the elements falls
into four blocks (s, p, d, f) in the periodic table.
z
z
Helium belongs to the s-subshell but is noted above
the p-subshell block on the periodic table.
• It does not share the same chemical & physical
properties of the elements in the s-subshells.
Hydrogen belongs to the s-subshell, but has properties
more similar to the elements in the p-subshell block.
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Electronic Structure and the
Periodic Table
Electronic Structure and the Periodic
Table
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Noble Gases (rare gases)
z
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All of them have completely filled electron shells.
Since they have similar electronic structures,
their chemical reactions are similar.
z
z
z
z
z
z
He
Ne
Ar
Kr
Xe
Rn
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1s2
[He] 2s2 2p6
[Ne] 3s2 3p6
[Ar] 4s2 4p6
[Kr] 5s2 5p6
[Xe] 6s2 6p6
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Representative Elements (Main Group)
z Are the elements in A groups on periodic chart.
These elements will have their “last” electron in an
outer s or p orbital.
These elements have fairly regular variations in
their properties.
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Electronic Structure and the Periodic Table
Electronic Structure and the Periodic Table
• d-Transition Elements
Group IA, IIA, IIIA
Æ form the main group metals
Group IVA, VA, VIA, VIIA Æ
form the main group nonmetals
Group VIIIA
Æ noble gases (inert and exist as
single atoms)
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– Elements on periodic chart
in B groups.
– Sometimes called transition
metals.
• Each metal has d electrons.
– ns (n-1)d configurations
• These elements make the
transition from metals to
nonmetals.
• Exhibit smaller variations from
row-to-row than the
representative elements.
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Electronic Structure and the Periodic Table
Electronic Structure and the Periodic
Table
• f - transition metals
– Sometimes called inner
transition metals.
• Electrons are added to f
orbitals.
• Electrons are added two shells
below the valence shell!
• Consequently, very slight
variations of properties from
one element to another.
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Outermost electrons have the greatest influence on the chemical
properties of elements.
The group number in the main groups indicates the number of
valence-shell electrons of elements in the group.
z Group IA elements have 1 electron in their valence-shell
1s1,2s1, 3s1, 4s1, 5s1, 6s1, 7s1
z Group IIA elements have 2 electrons in their valence-shell
1s2, 2s2, 3s2, 4s2, 5s2, 6s2, 7s2
z Group III A elements have 3 electrons in their valence-shell
ns2np1
z Group VA elements have 5 electrons in their valence-shell
ns2np3
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Periodic Properties of the Elements
Periodic Properties of the Elements
Atomic radii describes the relative sizes of atoms.
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The atomic radius of an element is defined as half the
distance between the nuclei of neighboring atoms
The periodic table is useful in determining
different properties of elements that:
z
z
If the element is nonmetal or metalloid; the
distance between nuclei joined by chemical
bond (covalent radius)
Help us understand and predict chemical bonding
Help us understand and predict certain trends associated
with the organization of the elements on the table
Atomic radii is determined by
measurement of the distance between the
nuclei of homonuclear molecules, e.g.,
Cl2 (198 x 10−12 m Æ 198 pm)
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p.236
Atomic Radii (Section 1.14)
Periodic Properties of the Elements
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Atomic radii increase within a column going
from the top to the bottom of the periodic table.
Atomic radii decrease within a row going from
left to right on the periodic table.
z
z
The atomic radii trend:
This last fact seems contrary to intuition.
How does nature make the elements smaller even
though the electron number is increasing?
Slide 7
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p.238
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Periodic Properties of the Elements
Figure 1.40 Effective Nuclear Charge
Atomic Radii
The reason the atomic radii decrease across a period is
due to shielding or screening effect.
z
z
Effective nuclear charge, Zeff, experienced by an electron is
less than the actual nuclear charge, Z.
The inner electrons block the nuclear charge’s effect on the
outer electrons by repulsion.
• Each electron, in an atom with many electrons, is
repelled by the other electrons.
• The outer electrons are shielded by the inner electrons
from the full attraction of the nucleus.
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Figure 12.38: Atomic
radii (in picometers)
for selected atoms.
Atomic Radii
The effective nuclear
charge increases
across the periods.
Consequently, the
outer electrons feel a
stronger effective
nuclear charge.
2s1
3+
Zeff decreases
• nuclear charge is the
positive charge of the
nucleus
• the size of an atom is
determined by the
effective nuclear
charge
• shielding electrons are
the filled energy levels
of an atom and they
tend to shield the outer
electrons from feeling
the full attraction of the
nucleus.
Zeff increases
1s2
Li has two shielding electrons (1s2)
The one valence electron (2s1) does
not feel the full effect of the 3+
nucleus
As electrons are added
to more energy levels,
the size of the atom
increases.
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Ionic Radius
Atomic Radii
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Arrange these elements based on their increasing
atomic radii.
z
Se, S, O, Te
z
P, Cl, S, Si
z
Ga, F, S, As
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• The ionic radius of an element is the distance between
the radii of the ions in an ionic solid (ionic crystal).
O < S < Se < Te
Cl < S < P < Si
F < S < As < Ga
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Ionic Radii
Ionic Radii
Cations (positive ions) are
always smaller than their
respective neutral atoms.
ƒ Cations (positive ions) are always smaller than their
respective neutral atoms.
Element
Li
Na
Atomic
Radius (Å)
1.52
1.86
Ion
Li+
Na+
Ionic
Radius (Å)
0.90
1.16
Anions (negative ions) are
always larger than their
respective neutral atoms.
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p.244
Ionic Radii
Ionic Radii
An isoelectronic series consist of ions that have the same
number of electrons.
• Anions (negative ions) are always larger than
their neutral atoms.
The Zeff has a more obvious effect on the metal ions.
Element
N
O
F
Atomic
Radius(Å)
0.75
0.73
0.72
Ion
N3−
O2−
F−
Ionic
Radius(Å)
1.71
1.26
1.19
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Ionic Radii
Ionic Radii
Note: This
trend is
generalized to
metal ions and
nonmetals ions
separately.
ƒ Cation (positive ions) radii decrease from left to
right across a period.
– Increasing nuclear charge attracts the electrons and
decreases the radius.
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Ion
Rb+
Sr2+
In3+
Ionic
Radii(Å)
1.66
1.32
0.94
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Ionic Radii
Ionic Radii
ƒ Anion (negative ions) radii decrease from left to
right across a period.
Arrange these elements based on their increasing
ionic radii.
– Increasing electron numbers in highly charged ions cause
the electrons to repel and increase the ionic radius.
z
Ga, K, Ca
z
Cl, Se, Br, S
Ga3+ < Ca2+ < K1+
Ion
N3-
O2-
F1-
Ionic
Radii(Å)
1.71
1.26
1.19
Cl1- < S2- < Br1- < Se2-
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Ionization Energy
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Ionization Energy
In an ionic compound (salt), the formation of the
electrostatic bond (ionic) is dependent upon the
ability to remove one or more electrons from
atoms.
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First ionization energy (IE1)
z
z
z
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The minimum amount of energy required to remove the
most loosely bound electron from an isolated gaseous atom
to form a 1+ ion.
This is the energy required to remove the highest-energy
electron.
eV = electronvolts for single atoms and kJ/mol for a
mole of atoms
Symbolically:
Atom(g) + energy (eV) → ion+(g) + eMg(g) + 738kJ/mol → Mg+ (g) + e-
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First Ionization Energies of Some Elements
Ionization Energy
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Elements with filled or half-filled subshells are stable and require
more energy to remove an electron than their neighbor to the right
in the period.
Second ionization energy (IE2)
The amount of energy required to remove the second electron
from a gaseous 1+ ion.
Symbolically:
z ion+ + eV → ion2+ + ez
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Mg+
+ 1451 kJ/mol →
Mg2+
+
Ne
2000
Ionization
Energy
(kJ/mol)
e-
Atoms can have 3rd (IE3), 4th (IE4), etc. ionization
energies.
He
2500
F
1000
Ar
N
1500
H
C
Be
Cl
O
B
500
Li
Na
0
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Si
Mg
1 2
3 4
5 6 7
Al
P
Ca
S
K
8 9 10 11 12 13 14 15 16 17 18 19 20
Atomic Number
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First Ionization Energies of Some Elements
Ionization Energies
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Ionization Energy
Ionization Energy
Periodic Trends for
Ionization Energy:
IE2 > IE1
It always takes more energy to
remove a second electron
from an ion than from a
neutral atom.
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•
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Arrange these elements based on their first
ionization energies.
z
Sr, Be, Ca, Mg
z
Al, Cl, Na, P
z
B, O, Be, N
Sr < Ca < Mg < Be
IE1 generally increases moving from IA elements to
VIIIA elements.
IE1 generally decreases moving down a group.
IE1 for Li > IE1 for Na, etc.
Na < Al < P < Cl
B < Be < O < N
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Electron Affinity (Ea )
Electron Affinity
Symbolically:
Electron affinity is the amount of energy change
when an electron is added to an isolated gaseous
atom.
| If an electron is added to an element, the amount
of energy change is exothermic, therefore, the
Ea is negative.
The more negative the energy change, the greater
the energy released when an electron is added.
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Atom (g) + e- → ion- (g)
If Ea is positive, energy must be added to
push an electron onto an atom.
„ If Ea is negative, energy is released when an
electron is added onto an atom
„
E.g. When an electron is added to a chlorine atom, the electron
has a lower energy.
The electron affinity of
E(Cl ) 
→ E(Cl – ) is negative
chlorine is negative.
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Electron Affinity
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Electron Affinity
Often, electron affinity is referred to as the energy
released.
z
z
Those elements which have a high electron affinity
release energy when an electron is added.
The general trend is that the electron affinity is greater
for the nonmetals, particularly the halogens, than for
metals.
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The Pauling Scale of Electronegativity
Electronegativity
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Electronegativity is a measure of the relative tendency
of an atom to attract electrons to itself when
chemically combined with another element.
z
z
z
Electronegativity is measured on the Pauling scale.
Fluorine is the most electronegative element.
Cesium and francium are the least electronegative elements.
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Electronegativity
Electronegativity
For the representative elements, electronegativity
usually increase from left to right across periods and
decrease from top to bottom within groups.
Arrange these elements based on their increasing
electronegativity.
z
Se, Ge, Br, As
z
Be, Mg, Ca, Ba
Ge < As < Se < Br
Ba < Ca < Mg < Be
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Figure 12.39: Special names for groups in the periodic table
The Properties of a Group
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Elements in a group tend to share similar
chemical properties with the other group
members.
Elements in the same group share the same
number of valence electrons and electronic
arrangement.
It is the number and type of valence electrons
that primarily determine an atom’s
chemistry.
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The Alkali Metals
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The Alkali Metals
The group IA metals are called the alkali metals and are the most
reactive metals.
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Group IA elements have low ionization
energies and will react readily with
nonmetals to form ionic solids.
Group IA elements are very good reducing
agents (they lose electrons and become
oxidized in ionic bond formation.)
Group IA elements react very exothermically
with water.
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End of Chapter 6
Periodic Trends
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It is important that you understand and know the
periodic trends described in the previous sections.
z
They will be used extensively in Chapter 7 to understand
and predict bonding patterns.
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