Transcript The History of the Modern Periodic Table

The Periodic
Table
During the nineteenth century,
chemists began to categorize the
elements according to similarities
in their physical and chemical
properties. The end result of
these studies was our modern
periodic table.
Dmitri Mendeleev
In 1869 he published a table of
the elements organized by
increasing atomic mass.
1834 - 1907
Lothar Meyer
At the same time, he published his own
table of the elements organized by
increasing atomic mass.
1830 - 1895
Henry Moseley
In 1913, through his work with X-rays, he
determined the actual nuclear charge
(atomic number) of the elements*. He
rearranged the elements in order of
increasing atomic number.
*“There is in the atom a fundamental
quantity which increases by regular
steps as we pass from each element to
the next. This quantity can only be the
charge on the central positive nucleus.”
1887 - 1915
Periodic Table
Geography
Periodic Table
• Columns called Families/Groups
– Family # indicates # valence (outer shell)
electrons
– Elements in same family have similar
properties
• Rows called Periods
– Row # indicates # energy levels in atom
The horizontal rows of the periodic table
are called PERIODS.
The elements in any group
of the periodic table have
similar physical and chemical
properties!
The vertical columns of the periodic table
are called GROUPS, or FAMILIES.
Metals/Nonmetals/Semiconductors
• Metals: excellent conductors of heat & electricity;
have luster, are ductile/malleable
• Nonmetals: poor conductors of heat & electricity;
are dull & brittle
• Semiconductors(Metalloids): elements that under
certain conditions conduct heat & electricity
Families of Elements
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Family 1:
Alkali Metals
Family 2:
Alkaline Earth Metals
Families 3 to 12:
Transition Metals
Family 13:
Boron Family
Family 14:
Carbon Family
Family 15:
Nitrogen Family
Family 16:
Oxygen Family
Family 17:
Halogens
Family 18:
Noble Gases
Three general groups: metals, nonmetals, &
semiconductors(metalloids)
Periodic Table
Periodic Law
When elements are arranged in order of
increasing atomic number, there is a
periodic pattern in their physical and
chemical properties.
Alkali Metals
Alkaline Earth Metals
Transition Metals
Metals
• Alkali metals (Family 1)
– Very reactive
– Has 1 valence electron
– When ionized has charge of 1+
• Alkaline Earth metals (Family 2)
– Reactive
– Has 2 valence electrons
– When ionized has charge of 2+
• Transition metals (Families 3 to 12)
– Somewhat reactive
– Valence electron number varies
– Ionized charge varies
These elements are also
called the rare-earth
elements.
InnerTransition Metals
Halogens
Noble Gases
Nonmetals
• Include H, some elements from families 13 to 16, all
elements from families 17 & 18. Zig-zag line divides metals
from nonmetals.
• Inert gases are unreactive; contain 8 valence electrons
• Halogens are very reactive; contain 7 valence
electrons; gain electrons becoming negatively
charged
• Elements in other families gain electrons to become
negatively charged
• These elements plentiful on Earth
Semiconductors
(aka Metalloids)
• Located along the zig-zag line
• Includes:
– Boron (B); Silicon (Si), Germanium (Ge), Arsenic (As), Antimony
(Sb), Tellurium (Te), Polonium (Po)
• Notice that Al is not considered a metalloid, it is
considered a metal
• Conduct heat & electricity under certain conditions
• B is hard & added to steel to increase hardness; Sb is
bluish-white and shin, Te is silvery-white & electrical
conductivity increases with light exposure, Si important in
solar cells & integrated circuits
Periodic Table Trends
#1. Atomic Size - Group trends
• As we increase the
atomic number (or
go down a group). .
.
• each atom has
another energy
level,
• so the atoms get
bigger.
H
Li
Na
K
Rb
#1. Atomic Size - Period Trends
• Going from left to right across a period, the size
gets smaller.
• Electrons are in the same energy level.
• But, there is more nuclear charge.
• Outermost electrons are pulled closer.
Na
Mg
Al
Si
P
S Cl Ar
#2. Trends in Ionization Energy
• Ionization energy is the amount of
energy required to completely remove
an electron (from a gaseous atom).
• Removing one electron makes a 1+
ion.
• The energy required to remove only
the first electron is called the first
ionization energy.
Ionization Energy
• The second ionization energy is the
energy required to remove the second
electron.
– Always greater than first IE.
• The third IE is the energy required to
remove a third electron.
– Greater than 1st or 2nd IE.
#3. Trends in Electronegativity
• Electronegativity is the tendency for an
atom to attract electrons to itself when it
is chemically combined with another
element.
• They share the electron, but how equally
do they share it?
• An element with a big electronegativity
means it pulls the electron towards itself
strongly!
Electronegativity Group Trend
• The further down a group, the
farther the electron is away
from the nucleus, plus the more
electrons an atom has.
• Thus, more willing to share.
• Low electronegativity.
Electronegativity Period Trend
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Metals are at the left of the table.
They let their electrons go easily
Thus, low electronegativity
At the right end are the nonmetals.
• They want more electrons.
• Try to take them away from others
• High electronegativity.
The periodic table is the most important
tool in the chemist’s toolbox!