Transcript Periodic Table Introduction

The Periodic Table
Early Organization
As early as the early Greeks, scientists
wanted to organize.
 They separated matter into Earth, Air, Fire
and Water.
 They even had more detail such as
combinations.
 For example: Earth and Fire = Lava
 Check out the game Little Alchemy on
GOOGLE – you might like it.

History of the Periodic Table
 In
the 1700s
scientists had
identified only 30
elements
 In the 1800s
there were 60
History of the Periodic Table
 Early
1800s
Dobereiner grouped
elements into triads
 Li, Na, K
History of the Periodic Table
 1865
J.A.R. Newlands
discovered properties
repeated themselves
every eighth element
 Called these octaves
Mendeleev
 1869
Dmitri Mendeleev
produced the first
periodic table for his
students
 Left blanks where
elements had yet to be
discovered
 Was organized by atomic
weight as that was the
standard at the time.
Here is one of the early official versions of
Mendeleev’s Periodic Chart
Notice it is in German, even though Mendeleev was Russian.
Properties of Germanium as Predicted by Mendeleev
Properties of Ekasilicon
Atomic weight
Density
Specific heat
Melting point
Oxide formula
Oxide density
Chloride formula
bp of chloride
Predicted in 1871
72
5.5 g/cm3
0.31 J/(°C · g)
Very high
RO2
4.7 g/cm3
RCl4
100°C
Properties of
Germanium
Atomic weight
Density
Specific heat
Melting point
Oxide formula
Oxide density
Cl-1 formula
bp of chloride
Predicted
in 1871
72
5.5 g/cm3
0.31 J/(°C · g)
Very high
RO2
4.7 g/cm3
RCl4
100°C
Observed in
1886
72.3
5.47 g/cm3
0.32 J/(°C · g)
960°C
GeO2
4.70 g/cm3
GeCl4
86°C
HenryMoseley

In 1913 Moseley assigned
elements atomic numbers
and rearranged periodic
table.
Table Terms
 Periodic
Law – when arranged by
increasing atomic number elements
repeat similar chemical and physical
properties
 Groups or Families are the columns
on the periodic table
 Periods are the rows going across.
Major Groups on the Periodic Table
 Oxygen Group
Metals
(Chalcogens)
 Alkaline Earth
 Halogens
Metals
 Transition Metals  Noble Gases
 Actinide Series
 Boron Group
 Lanthanide
 Carbon Group
Series
 Nitrogen Group
 Alkali
6.1
Metals, Nonmetals, and Metalloids
Check out Theodore Gray’s App – The elements
6.1
Metals, Nonmetals, and
Metalloids
6.1
Metals, Nonmetals, and
Metalloids
Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
Metals
Good Conductors of
heat and electricity
 Luster – Shiny
 Malleable – pounded
into thin sheets
 Ductile – pulled into a
wire
 Mercury is the only
liquid metal at room
temp

Non-Metals
 Most
are gases
 Solids are brittle (S &
P)
 Bromine is the only
liquid nonmetal at
room temp
Metalloids
 Properties
of metals
and non-metals
 Semi conductors
 Make very good
computer chips
Atomic Radius

Atomic radius is the distance from the
atom’s nucleus to its outer edge.
– In the same energy level, more protons
exert a stronger pull towards the nucleus
Showing the Trend of Atomic Size
Which has larger atomic radius?
 Na
or Rb
 P or Cl
You can tell by looking at
the chart and knowing
the trend.
Ionization Energy

Ionization energy is the energy needed
to remove one electron
 Na (g)  Na+1 (g) + 1 e-1
 Metals are more likely to give up an
electron than nonmetals.
Trends in Ionization Energy
It is good to know that Fluorine has the highest Ionization Energy
Which has larger ionization energy?
H
or Cs
 Li or N
Ionic Radius
 Ionic
radius is the distance from the
ion’s nucleus to its outer edge.
 Non-metal ions get larger with a
negative charge
 Metal ions get smaller with a positive
charge
 This is because more protons are
pulling on fewer electrons
Anions are going to be larger than the atom. Cations are
going to be smaller than the atom.
Trends in Ionic Size
 Relative Sizes of Some Atoms and Ions
Which of the following is larger?
O
or O-2
 K or K+
Any Questions????
Electronegativity
 Electronegativity
reflects an atom’s
ability to attract electrons
 Cs & Fr have the lowest
electronegativities; F has the highest
Reactivity of Alkali and Alkaline
Earth Metals
 Metals
become more
reactive as you
move down the
group
 Metals become less
active when moving
left to right
Why Are There Patterns ?
 Elements
have physical and chemical
properties based upon their valence
electrons.
 Valence electrons are the electrons in
the outer most energy level (s & p
orbitals).
 The number of valence electrons may be
determined by using the periodic table.
Why Are There Patterns?
 When
you look at an atom you are
observing the valence electrons
 Duet rule - only 2 electrons fill the first
energy level
 Octet rule – 8 valence electrons is
considered to be a full set
Electron Affinity

Electron affinity is the energy change that
occurs when an atom gains an electron.
 F (g) + e-1  F-1 (g)
 Most atoms give off energy when an electron
is gained (negative).
 Nonmetals have more of an electron affinity
than metals.
 EA decreases when moving down the group.
 EA increases moving from left to right in a
period.