Download Periodic Trends

Chapter 6
The Periodic Law
http://www.privatehand.com/flash/elements.html
History of the Periodic Table
1869 – Dmitri Mendeleev published his
periodic table. He had arranged it by
grouping together the elements that had
similar properties and
by increasing atomic
masses. His periodic
table left empty spaces
for new elements that
would be discovered.
Mendeleev’s List of elements in Russian Circa 1869
Periodic Table in English (Circa 1891)
Periodic Table circa 1898
History of the Periodic Table

1911 – Henry Moseley (a student of Ernest
Rutherford) rearranged a few elements on
the periodic table so that elements were
arranged by increasing atomic
number rather than by atomic
mass.
History of the Periodic Table

1944 – Glenn T. Seaborg rearranged the
periodic table to make it look like it does
today. He moved the Actinide Series and
the Lanthanide Series
elements to the bottom
of the periodic table.
http://livingtextbook.or
egonstate.edu/media/vi
d/lbl5a3.mov
Periodic Table Circa 1944
Modern Periodic Table
Parts of the Periodic Table
The
periodic table can be
divided and labeled using
several methods.
Parts of the Periodic Table
Elements are arranged:
Vertically columns are
called Groups
Horizontal rows are
called Periods
Periodic Families
Alkali
Metals
Alkaline Earth Metals
Halogens
Noble gases
Parts of the Periodic Table
Metals
Non-metals
Metalloids
Parts of the Periodic Table
Main
Group or
Representative Elements
Transition Metals
Rare earth elements
Trans Uranium Elements
The Periodic Law
The
physical and chemical
properties of the elements are
periodic functions of their
atomic numbers
Periodic Trends
If you understand the trends on the periodic
table, you can predict almost anything about
any element on the periodic table.
We will study:
• Atomic Radii
• Ionic Radii
• Valence Electrons
• Reactivity
• Electronegativity
• Electron Configuration
Atomic Radius

Define – One-half the distance between the
nuclei of identical atoms that are bonded
together.
Periodic Trend: The
atomic radius
decreases as you go
across the periodic
table and increases as
you go down the
periodic table
Periodic Trends
 As
you go across a period, the radius gets
smaller.
 Electrons are in same energy level.
 More nuclear charge.
 Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
Group trends
 As
we go down a
group...
 each atom has
another energy
level,
 so the atoms get
bigger.
H
Li
Na
K
Rb
Trends in Ionic Size
 Cations
form by losing electrons.
 Cations are smaller than the atom they
come from.
 Metals form cations.
 Cations of representative elements
have noble gas configuration.
Ionic size
 Anions
form by gaining electrons.
 Anions are bigger than the atom they
come from.
 Nonmetals form anions.
 Anions of ‘main’ groups elements have
noble gas configuration.
Group trends
Adding energy level
 Ions get bigger as you
go down.

Li1+
Na1+
K1+
Rb1+
Cs1+
Periodic Trends
 Across
the period, nuclear charge
increases so they get smaller.
 Energy level changes between anions
and cations.
32N
Li1+
B3+
Be2+
C4+
O
F1-
Valence Electrons
Define:
The electrons available
to be lost, gained, or shared in
the formation of compounds.
The
electrons in the highest
energy level
Valence Electrons
Periodic Trends:
Group 1 = 1 valence electron = 1+ Oxidation Number
Group 2 = 2 valence electrons = 2+ Oxidation Number
Group 13 = 3 valence electrons = 3+ Oxidation Number
Group 14 = 4 valence electrons = 4+/4- Oxidation Number
Group 15 = 5 valence electrons = 3- Oxidation Number
Group 16 = 6 valence electrons = 2- Oxidation Number
Group 17 = 7 valence electrons = 1- Oxidation Number
Group 18 = 8 valence electrons = 0 Oxidation Number
0
1+
3+
4+
4-
3-
2-
1-
He
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Ga
Ge
As
Se
Br
Kr
Rb
Sr
In
Sn
Sb
Te
I
Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
Fr
Ra
H
2+
Li
Reactivity
Reactivity increases as you go down the
columns of metallic elements.
 Reactivity decreases as you go down the
columns of non-metallic elements.
 Watch the video to see what that means.

Electronegativity
The tendency for an atom to attract
electrons to itself when it is chemically
combined with another element.
 High electronegativity means it pulls the
electron toward it.

Group Trend
 The
further down a group, the farther
the electron is away, and the more
electrons an atom has.
 More willing to share.
 Low electronegativity.
Periodic Trend
 Metals
are at the left of the table.
 They let their electrons go easily
 Low electronegativity
 At the right end are the nonmetals.
 They want more electrons.
 Try to take them away from others
 High electronegativity.
Electronegativity
(important to determine bond type)
Electron Configuration and the
Periodic Table
Electron Configuration
of Main Group Elements
Group
1
2
13
15
17
18
Period #(+)
s1
s2
s2p1
s2p3
s2p5
s2p6
Example
Na = 3s1
Ba = 6s2
Ga = 4s24p1
Sb = 5s25p3
2 5
Br = 4s 4p
Rn = 6s26p6
Electron Configuration
Transition Elements
Period # s2 + Period # (-1) d1 – 10
Examples: Sc = 4s23d1
Zn = 4s23d10
Mo = 5s24d4
Ir = 6s25d7
Electron Configuration for
Lanthanide and Actinide Series
Period # s2 + Period # (- 2) f 1 - 14
Examples: Ce = 6s2 4f1
Ho = 6s24f10
U = 7s25f3
Bk = 7s25f8
Atomic size increases,
shielding constant
Ionic size increases
Ionization energy, Electronegativity,
and Electron Affinity INCREASE