Transcript The Periodic Table - Jackson County School District

The Periodic Table
Dmitri Mendeleev
(1834 – 1907)
• He organized elements into
the first periodic table
• He arranged elements by
increasing atomic mass
Henry Moseley
(1913)
• He arranged elements
according to atomic number
rather than atomic mass
• The modern periodic table is
arranged by atomic number
Periodic Law
The periodic law states that
there is periodic repetition
of chemical and physical
properties of elements
The Modern Periodic Table
There are 18 groups (columns
up and down)
The group A’s (the tall columns)
are called representative
elements
• The group B’s (the middle
columns) are called transition
metals
There are seven periods
(rows across the periodic
table)
Metals are to the LEFT of the
zig-zag line (except hydrogen!)
Metals in yellow
Nonmetals are to the RIGHT of
the zig-zag line
nonmetals in red
Metalloids
• Metalloids are those
elements ON the zig-zag line
Metalloids border the
zig-zag line
Now . . .YOU fill in the
chart using your book!
Metals
• solid at room temperature
• shiny (have luster) and smooth
• good conductors of heat and
electricity
Metals
• malleable – “bendable” (can
be pounded into sheets)
• ductile - can be pulled
into wires
Metals
• react with acids
• mercury (Hg) is the only
LIQUID metal
Nonmetals
• generally gases or brittle, dull
looking solids at room
temperature
• poor conductors of heat and
electricity
• Bromine (Br) is the only
LIQUID nonmetal
Metalloids
• sometimes called semimetals
• metalloids have properties
of both metals and
nonmetals
Metalloids
• silicon and germanium are two
of the most important
metalloids (they’re used in
computer chips and solar cells)
Trends of the
Periodic Table
Periodic Law
If elements are organized
according to atomic number,
their properties will repeat
periodically
Four Periodic Trends:
1.Atomic radii
2.Ionic radii
3.Electronegativity
4.Ionization energy
Atomic Radius
The atomic radius basically
tells you the size of the
atom. It is half the distance
between two nuclei of identical
atoms bonded together.
radius
The trend: atomic radii
DECREASE across a period
• Why?
• Each time a positive proton is
added to the nucleus, the
negative electrons feel a greater
attraction to the positivelycharged nucleus and get “pulled
in” tighter
Decreasing – getting smaller!
Trend: atomic radii
INCREASE down a group
• Why?
• electrons are added to
higher and higher energy
levels as you go down
Atomic radii DECREASE
down a group!
The farther the electrons
from the nucleus, the larger
the atomic radii!!!!!
Try these . . .
1. Which element has the
larger atomic radius: C or F?
 carbon
2.Which element has the smaller
atomic radius: Ar or Kr?
 argon
Ionic Radii
• Is basically the size of an ion
or half the distance between
the nuclei of two ions bonded
together
What is ion???
Ionic Radii
• Ion – an atom with a
charge (+ or - )
• An ion is formed when
atoms lose or gain electrons
• What happens to an atom if
it LOSES an electron?
 it loses a negative charge
so it becomes POSITIVE
Na
+1
• A positively charged ion is
called a cation
Na
+1
Positive ions (cations) are
smaller than the atoms they
come from because they lose
electrons making the atom
smaller.
Na
Na
+1
• What happens to an atom if
it GAINS an electron?
 It gets more negative
(so it has a negative charge)
Cl
-1
A negatively charged ion is
called an anion.
Cl
-1
Negative ions (anions) are
LARGER than the atoms they
come from because they gain
electrons – making the atom
LARGER!
Cl
Cl
-1
The trend:
• See the board
Remember . . .
• all atoms want a full octet
(8 valence electrons)
• atoms with 1 valence electron
will give up that electron
VERY QUICKLY to become
stable
example:
• sodium has one valence
2
2
6
1
electron: 1s 2s 2p 3s
• if sodium gives it away,
then the configuration will be:
1s22s22p6
• sodium will have a full octet
• atoms with 7 valence
electrons will hold on to those
electrons VERY TIGHTLY
• they try to get one more and
become stable
Ionization Energy
• The amount of energy needed
to remove an electron
• think of it as: how tightly an
atom holds on to its electrons
The trend:
• ionization energy INCREASES
across a period
• Why?
• the more valence electrons
an element has, the more
difficult it is to remove
them!
The trend:
• Ionization energy DECREASES
down a group
• Valence electrons in higher
energy levels are NOT held
as tightly because they are
farther from the nucleus
• Therefore, it is easier to
remove an electron that is
farther from the nucleus
Try these . . .
• Which has a higher
ionization energy: Na or Cl
 Chlorine
• Which has a lower ionization
energy: Li or O
 Lithium
Electronegativity
• The ability of an atom to
attract electrons to itself
The Trend: electronegativity
INCREASES across a period
• Why?
• atoms are trying harder to
attract electrons to get a full
octet
The trend: electronegativity
DECREASES down a group
• Why?
• it is harder to hold on to the
electrons that are farther
away from the nucleus
Try these . . .
• Which element is
electronegative?
 Fluorine
• Which element is
electronegative?
 Boron
more
F or Br
more
B or Ca
Finished!