Download Periodic Table Notes Ch. 6 ppt

Chapter 6
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 Draw
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
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Chapter 6.1-6.2
 Periodic Law
 Group
 Period
 Representative
Element
 Transition Element
 Metal
 Alkali Metal
 Alkaline Earth Metal
 Transition
Metal
 Inner Transition
Metal
 Lanthanide Series
 Actinide Series
 Nonmetal
 Halogen
 Noble Gas
 Metalloid
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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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Properties of Germanium (Ge)
Property
Atomic Mass
Predicted (1869)
72 u
Color
Density
Melting Point
Dark gray
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
EsCl4
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 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.
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Columns,
Groups or
Alkali Metals
Families
Alkaline Earth Metals
Representative Elements
Periodic Table
Noble
Halogens Gases
Transition Elements
Periods
Inner
Transition
Elements
Metals
Metalloids
Nonmetals
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 What
are some of the elemental
properties that make the periodic table,
well, periodic?
 Classification by metals, nonmetals and
metalloids
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

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 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
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 Transition
B Group)
 d-block
Elements (Sometimes called
elements have 1-10 electrons in d-
orbitals
 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
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 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
 Ionization energy
 Electronegativity
 Metallic Character
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
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 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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 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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Orbital energies “drop” as Z* increases
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 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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 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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Radius (pm)
250
K
1st transition
series
3rd period
200
Na
2nd period
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
30
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Atomic Number
25
40
 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
2p
2p
2s
2s
1s
1s
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Atoms tend to gain, lose, or share
electrons to get
8 valence electrons
(except small atoms up to Boron)
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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?
2. Write the electron configuration and orbital
box diagram for O when it is an ion.
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+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming a
positive
ion.
 Positive
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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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.)

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See Figure 6-14
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Why do metals lose
electrons in their
reactions?
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take
on electrons?
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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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IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ  Mg+ (g) + e-
Mg+ (g) + 1451 kJ  Mg2+ (g) + eMg+ has 12 protons and only 11 electrons.
Therefore, IE for Mg+ > Mg.
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1st: Mg (g) + 735 kJ  Mg+ (g) + e2nd: Mg+ (g) + 1451 kJ  Mg2+ (g) + e-
3rd: Mg2+ (g) + 7733 kJ  Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n.
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1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
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Atomic Number
38
33
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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.
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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 IE
decreases down a group
 Because size increases.
 Ability to lose electrons
generally increases down
the periodic table.
 See reactions of Li, Na, K
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 Which
element in each pair has the
larger 1st ionization energy?
A. Na or Al
B. Ar or Xe
C. Ba or Mg
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Lithium
Sodium
Potassium
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*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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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?
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 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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