Chapter 6 - The Periodic Table

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Chapter 6 - The Periodic Table
 Draw
Chapter 6
the periodic table and label
the electron blocks and areas of
non-metals, metals, and metalloids.
 Relate the Lewis dot structure to its
place in the periodic table.
 Explain periodic trends as one moves
along periods and down groups in
the periodic table
1
Chapter 6.1-6.2
 Periodic Law
 Group
 Period
 Representative
Element
 Transition Element
 Metal
 Alkali Metal
 Alkaline Earth Metal
 Transition
 Inner
Metal
Metal
Transition
 Lanthanide
Series
Series
 Nonmetal
 Halogen
 Noble Gas
 Metalloid
 Actinide
2
Dmitri Mendeleev
noticed in his table
that there were
repetitions of
physical and
chemical properties
when the elements
were arranged by
atomic mass.
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4
Properties of Germanium (Ge)
Property
Predicted (1869)
Atomic Mass
72 u
Color
Dark gray
Density
Melting Point
5.5 g/mL
High
Density of Oxide
4.7 g/mL
Actual (1886)
Oxide solubility in Slightly dissolved
HCl
by HCl
Formula of chloride
5
EsCl4
6
1
Chapter 6 - The Periodic Table
 Periodic
Law states that chemical
and physical properties repeat in
regular cyclic patterns when they
are arranged by increasing atomic
number.
 Starts
with metals at left and goes to
non-metal (noble gas) on right
 Properties change in orderly progression
across a period.
7
Columns,
Groups or
Periodic Table
Alkali Metals
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Noble
Halogens Gases
Families
Alkaline Earth Metals
Representative Elements
 What
are some of the elemental
properties that make the periodic table,
well, periodic?
 Classification by metals, nonmetals and
metalloids
Transition Elements
Periods
Metals - shiny ductile, malleable solids, good
conductors of heat and electricity
 Nonmetals - dull, brittle solids; or gas, poor
conductors of heat and electricity
 Metalloids - have chemical and physical
properties of both metals and nonmetals

Inner
Transition
Elements
Metals
Metalloids
Nonmetals
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10
 Representative
called A Group)
Elements (Sometimes
 Group # = number of valence electrons
 Means similar Lewis dot structure and
similar properties.
 s-block elements have 1-2 electrons in
s-orbital
 p-block elements have 1-6 electrons in
p-orbitals
 Noble gases have filled valence shells
 Energy
level of valence electrons is at
energy level given by period (row)
number
11
 Transition
B Group)
Elements (Sometimes called
 d-block
elements have 1-10 electrons in dorbitals
 Columns 3-12 in periodic table
 Energy
level of valence electrons at n and
partially filled n-1 d orbitals (example: 4s
and 3d)
 f-block
(Lanthanides and Actinides) have
1-14 electrons in f-orbitals
12
2
Chapter 6 - The Periodic Table
 Fill
in the missing info for the following elements:
Configuration
Group
Period
7 (7B)
4
Block
[Ne]3s2
[He]2s1
[Kr]5s24d105p5
 Identify
the element fitting the description.
a) Group 2 (2A) element in 4th period:
b) Noble gas in 5th period:
c) Group 12 (2B) element in 4th period:
d) Group 16 (6A) element in 2nd period:
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 Effective
Nuclear Charge (Z*) – Not in book!
 Shielding (Not in book)
 Ion
 Ionization Energy
 Octet Rule
 Metallic Character (Not in book)
 Electronegativity
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 Atomic
and ionic size
energy
 Electronegativity
 Metallic Character
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
 Ionization
15
 Z*
is the nuclear charge experienced by the
outermost electrons. (Note: not in book!)
Z* increases across a period owing to shielding by
inner electrons.
 Shielding is blocking by inner electrons.

For a period (row), the number of shielding electrons
remain the same, but the number of protons in the
nucleus increases.
 Example: All elements in the second period have the
same underlying [He] noble gas configuration.
However, the number of protons increase from left to
right.

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16
 So
we can estimate as
Z* = [ Z - (no. inner electrons) ]
or
Z* = Z – S (inner electrons)
 Z is total number of electrons
 S is the number of electrons blocking the valence
shell electrons, the underlying noble gas electrons.
 Charge felt by 2s e- in Li
Z* = 3 - 2 = 1
 Be
Z* = 4 - 2 = 2
B
Z* = 5 - 2 = 3
and so on!
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3
Chapter 6 - The Periodic Table
Orbital energies “drop” as Z* increases
 Atomic
size is a periodic trend influenced by
electron configuration.
 For
metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the
element.
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20
 For
other elements, the atomic radius is half
the distance between nuclei of identical
atoms that are bonded together.
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 Size
(radius) goes UP on going
down a group. See previous slide.
 Because electrons are further
from the nucleus, there is less
attraction.
 Size (radius) goes DOWN on
going across a period.
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Size (radius) decreases across a period owing
to increase in Z*. Each added electron feels a
greater and greater positive charge.
Note: Electrons in the same energy level don’t
shield each other too much.
Large
Small
Increase in Z*
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4
Chapter 6 - The Periodic Table
 The
radius of an atom when it has
become an ion.
 An ion is an atom or bonded group of
atoms that has an overall positive or
negative charge.
 An atom acquires a positive charge by
losing electrons or negative charge by
gaining electrons!!
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To form positive ions from elements remove 1
or more e- from subshell of highest n [or
highest (n + l)].
Al: [Ne] 3s2 3p1 - 3e-  Al3+: [Ne] 3s0 3p0
3p
3p
3s
3s
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Atoms tend to gain, lose, or share
electrons to get
8 valence electrons
(except small atoms up to Boron)
2p
2p
2s
2s
1s
1s
27
1. Write the electron configuration and orbital
box diagram for Mg when it is an ion. Hints:
What is its noble gas configuration? What will
they do to get an octet?
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+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming a
positive
ion.
 Positive
2. Write the electron configuration and orbital
box diagram for O when it is an ion.
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ions are SMALLER than the atoms
from which they come.
 The electron/proton attraction has gone
UP and so size DECREASES.
 Electron Configuration as ion is: [He] 2s0
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5
Chapter 6 - The Periodic Table
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming a
negative
ion.
Negative ions are LARGER than the atoms from
which they come.
 The electron/proton attraction has gone
DOWN and so size INCREASES.
 Trends in ion sizes are the same as atom sizes.
 Electron configuration as ion: 1s22s22p6 (just
like neon.)

See Figure 6-14
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Why do metals lose
electrons in their
reactions?
32
IE = energy required to remove an electron
from an atom in the gas phase.
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take
on electrons?
33
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ  Mg+ (g) + e-
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1st: Mg (g) + 735 kJ  Mg+ (g) + e2nd: Mg+ (g) + 1451 kJ  Mg2+ (g) + e-
Mg (g) + 738 kJ  Mg+ (g) + e-
3rd: Mg2+ (g) + 7733 kJ  Mg3+ (g) + eMg+ (g) + 1451 kJ  Mg2+ (g) + eMg+ has 12 protons and only 11 electrons.
Therefore, IE for Mg+ > Mg.
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Energy cost is very high to dip into a
shell of lower n.
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6
Chapter 6 - The Periodic Table
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38
As Z* increases, orbital energies
“drop” and IE increases.
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 IE
increases across a period
because Z* increases.
 Metals lose electrons more
easily than nonmetals.
 Nonmetals lose electrons with
difficulty.
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 IE
decreases down a group
size increases.
 Ability to lose electrons
generally increases down
the periodic table.
 See reactions of Li, Na, K
 Because
High ionization energy: atoms want
to hold on to electrons; likely to form
negative ion
Low ionization energy: atom gives up
electron easily; likely to form positive
ion
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42
7
Chapter 6 - The Periodic Table
 Which
element in each pair has the
larger 1st ionization energy?
A. Na or Al
B. Ar or Xe
C. Ba or Mg
Lithium
Sodium
43
Potassium
44
*Note: ‘metallic character’ not in book.
An element with metallic character is one
that loses electrons easily.
Metallic character:
• is more prevalent in metals on left side of
periodic table
• is less for nonmetals on right side of
periodic table that do not lose electrons
easily
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 Relative ability
of an element to attract
electrons in a chemical bond.
Ionization energy reflects ability of atom to
attract electrons in an isolated atom
 Generally, the higher the ionization energy of an
atom, the more electronegative the atom will be
in a molecule

 There
are many electro negativity scales –
we’ll use the one by Linus Pauling (values
dimensionless)
 Will be used to determine things like
polarity of a chemical bond.
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8
Chapter 6 - The Periodic Table
Decreases
 Why?
down a group
Due to greater atomic radius
Increases
across a period
 Why?
Increased positive charge in
nucleus (Greater Z*)
Same
trend as for ionization
energy. Surprised?
 Moving Left  Right (periods)
 Z* Increases
 Atomic & ionic Radius Decrease
 Ionization Energy Increases
 Electronegativity Increases
 Metallic Character Decreases
 Moving Top  Bottom (groups)
 Z* is roughly constant, but val e- distance
increases
 Atomic & Ionic Radius Increase
 Ionization Energy Decreases
 Electronegativity Decreases
 Metallic Character Increases
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a) Electronegativity
b) Ionic Radius
c) Atomic Radius
d) Ionization Energy
e) Metallic character
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