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Interpreting the Periodic Table
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Interpreting the Periodic Table
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General Properties of Metallic Elements:
1.
Typically they have a shiny luster.
2.
Relatively high density.
3.
Malleable ( they can be hammered into thin sheets).
4.
Ductile (they can be drawn into a wire).
5.
Good conductors of electricity.
6.
Reflect light and heat.
7.
High melting and boiling points so they are solids at
room temperature (except Hg).
8.
Lose electron(s) forming positive ions (cations).
9.
Combine with non-metals.
10. Do not readily combine with each other.
Examples of Metals
Potassium, K
reacts with
water and
must be
stored in
kerosene
Copper, Cu, is a relatively soft
metal, and a very good electrical
conductor.
Zinc, Zn, is more
stable than
potassium
Mercury, Hg, is the only
metal that exists as a liquid
at room temperature
11 p
11 e
11 p
10 e
• Metals lose electrons forming positive ions.
• The radius of the cation is always smaller than the
radius of the parent atom.
General Properties of Non-metallic Elements:
1.
Poor conductors of heat and electricity.
2.
Not malleable or ductile; fragile.
3.
Low densities.
4.
Low melting and boiling points so they can be
gases, liquids, or solids.
5.
Gain electron(s) forming negative ions (anions).
6.
Combine with metals.
7.
Combine with each other to a limited extent.
17 p
17 e
17 p
18 e
• Non metals gain electrons forming negative ions.
• The radius of the anion is always larger than the
radius of the parent atom.
Metalloids: elements that lie in the
colored stair step line of the Periodic
Table. They have both metallic and
non-metallic properties.
These elements are weak conductors of electricity, which makes
them useful semiconductors in the integrated circuits of
computers.
Physical state of elements
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Chemical Periodicity
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Periodic Properties
of the Elements
Atomic Radii
Atomic radii describes the
relative sizes of atoms. It is
understood as the distance from
the nucleus to the outermost
occupied energy level.
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.
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Atomic Radii
Example: Arrange these elements based on
their atomic radii.

Se, S, O, Te
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Atomic Radii
Example: Arrange these elements based on
their atomic radii.

P, Cl, S, Si
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Atomic Radii
Example: Arrange these elements based on
their atomic radii.

Ga, F, S, As
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Ionization Energy
First ionization energy (IE1)

The minimum amount of energy required to remove the
most loosely bound electron from an isolated gaseous
atom to form a 1+ ion.
Symbolically:
Atom(g) + energy  ion+(g) + e-
Mg(g) + 738kJ/mol  Mg+ + e-
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Ionization Energy
Second ionization energy (IE2)

The amount of energy required to remove the
second electron from a gaseous 1+ ion.
Symbolically:

ion+ + energy  ion2+ + eMg+ + 1451 kJ/mol Mg2+ + e-
•Atoms can have 3rd (IE3), 4th (IE4), etc.
ionization energies.
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Ionization Energy
Periodic trends for Ionization
Energy:
1.
IE2 > IE1
It always takes more
energy to remove a second
electron from an ion than
from a neutral atom.
2.
IE1 generally increases moving
from left to right in the same
period.
3.
IE1 generally decreases
moving down a group.
IE1 for Li > IE1 for Na, etc.
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Ionization Energy
Example: Arrange these elements based on
their first ionization energies.

Sr, Be, Ca, Mg
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Ionization Energy
Example: Arrange these elements based on
their first ionization energies.

Al, Cl, Na, P
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Ionization Energy
Example: Arrange these elements based on
their first ionization energies.

B, O, Be, N
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Ionization Energy
Group
and
element
IE1
(kJ/mol)
IE2
(kJ/mol)
IE3
(kJ/mol)
IE4
(kJ/mol)
IA
Na
IIA
Mg
IIIA
Al
IVA
Si
496
738
578
786
4562
1451
1817
1577
6912
7733
2745
3232
9540
10,550
11,580
4356
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Ionization Energy
The reason Na forms Na+ and not Na2+ is
that the energy difference between IE1 and
IE2 is so large.

Requires more than 9 times more energy to
remove the second electron than the first one.
The same trend is persistent throughout the
series.
Thus Mg forms Mg2+ and not Mg3+.
 Al forms Al3+.

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Ionization Energy
Example: What charge ion would be expected for
an element that has these ionization energies?
IE1 (kJ/mol)
IE2 (kJ/mol)
IE3 (kJ/mol)
IE4 (kJ/mol)
IE5 (kJ/mol)
IE6 (kJ/mol)
IE7 (kJ/mol)
IE8 (kJ/mol)
1680
3370
6050
8410
11020
15160
17870
92040
Notice that the
largest increase
in ionization
energies occurs
between IE7 and
IE8. Thus this
element would
form a 1- ion.
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Ionization Energy
Example: What charge ion would be expected for
an element that has these ionization energies?
IE1 (kJ/mol)
IE2 (kJ/mol)
IE3 (kJ/mol)
IE4 (kJ/mol)
IE5 (kJ/mol)
IE6 (kJ/mol)
IE7 (kJ/mol)
IE8 (kJ/mol)
1680
3370
11586
12410
13020
15160
17870
18040
Notice that the
largest increase
in ionization
energies occurs
between IE2 and
IE3. Thus this
element would
form a 2+ ion.
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Ionization Energy
Example: What charge ion would be expected for
an element that has these ionization energies?
IE1 (kJ/mol)
IE2 (kJ/mol)
IE3 (kJ/mol)
IE4 (kJ/mol)
IE5 (kJ/mol)
IE6 (kJ/mol)
IE7 (kJ/mol)
IE8 (kJ/mol)
1680
3370
4586
5410
6020
7160
17870
18040
Notice that the
largest increase
in ionization
energies occurs
between IE6 and
IE7. Thus this
element would
form a 2- ion.
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Ionic Radii
Cations (positive ions) are always smaller than
their respective neutral atoms.
Element
Li
Be
Atomic
Radius (Å)
Ion
1.52
1.12
Li+
Be2+
Ionic
Radius (Å)
0.90
0.59
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Ionic Radii
Cations (positive ions) are always smaller than
their respective neutral atoms.
Element
Na
Mg
Al
Atomic
Radius (Å)
1.86
1.60
1.43
Ion
Na+
Mg2+
Al3+
Ionic
Radius (Å)
1.16
0.85
0.68
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Ionic Radii
Anions (negative ions) are always larger
than their neutral atoms.
Element
N
O
F
Atomic
Radius(Å)
Ion
0.75
0.73
0.72
N3-
O2-
F1-
Ionic
Radius(Å)
1.71
1.26
1.19
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Ionic Radii
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Ionic Radii
Cation (positive ions) radii decrease from left to
right across a period.

Increasing nuclear charge attracts the electrons and
decreases the radius.
Ion
Rb+
Sr2+
In3+
Ionic
Radii(Å)
1.66
1.32
0.94
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Ionic Radii
Anion (negative ions) radii decrease from left to
right across a period.

Increasing electron numbers in highly charged ions
cause the electrons to repel and increase the ionic
radius.
Ion
N3-
O2-
F1-
Ionic
Radii(Å)
1.71
1.26
1.19
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Ionic Radii
Example: Arrange these elements based on
their ionic radii.

Ga, K, Ca
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Ionic Radii
Example: Arrange these elements based on
their ionic radii.

Cl, Se, Br, S
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Electronegativity
Electronegativity is a measure of the relative
tendency of an atom to attract electrons to itself
when chemically combined with another element.
 Electronegativity is measured on the Pauling
scale.
 Fluorine is the most electronegative element.
 Cesium and francium are the least
electronegative elements.
For the representative elements, electronegativities
usually increase from left to right across periods
and decrease from top to bottom within groups.
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Electronegativity
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Electronegativity
Example: Arrange these elements based on
their electronegativity.

Se, Ge, Br, As
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Electronegativity
Example: Arrange these elements based on
their electronegativity.

Be, Mg, Ca, Ba
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See animation for Summary of
Trends in the Periodic Table
http://www.learnerstv.com/animation/animation.php?ani=56&cat=chemistry
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