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Periodic Table
PERIODIC LAW
PERIODS AND GROUPS
IONIZATION ENERGY
ELECTRONEGATIVITY
ATOMIC RADIUS
Periodic Law
“When elements are arranged
in order of increasing atomic
mass, periodic, or repeating,
trends occur.”
In 1869, Mendeleev arranged
the elements in order of mass
he and his students studied
according to trends such as
conductivity, ductility,
malleability, etc.
Now the table is in atomic
number order!
One Such Trend
Metals
Non-metals
 On the left side of the
 On the right side of the







periodic table
Great conductors of
electricity and heat
Ductile
Malleable
Lusterous
Solids (except Hg)
Lose e- when bonding




periodic table
Not conductors of
electricity and heat
Brittle
Dull
Mostly gases (Br is a
liquid)
Tend to gain e- when
bonding
Metalloids
 Metalloids have properties of both metals and
nonmetals such as poor conductors, may be shiny,
etc.
 Metalloids are in the middle of the periodic table and
only include elements in which a whole side touched
the staircase.
Periods
 Periods represent horizontal rows on the periodic
table and are numbered 1-7. These numbers can also
represent the number of energy levels.
Groups
 Groups represent vertical columns on the table and
are numbered 1-18 (different on some other tables).
Elements in the same group have the same number
of valence electrons and similar properties.
 Let’s examine some groups…
Alkali Metals: Group 1
 Very reactive (not found
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



in elemental form)
One valence e-: 1+ charge
soft
Metallic character
increases down the group
mp and bp decreases
down the group but all
very high
Strong exothermic
reaction with water
Alkaline Earth Metals: Group 2
 Reactive (not as
much as alkali
metals)
 Two valence e 2+ charge
 High mp and bp
Transition Metals
 Groups 3-12
 Form colored
compounds and
solutions
 Various charges
Colloidal Silver’s effects on people
Halogens: Group 17
 Form diatomic
molecules
 Mostly 1- charge,
though others
possible
 7 valence e Very reactive
 Low mp and bp
Noble Gases: Group 18
 Also known as “inert




gases” because they do
not react
Zero charge
8 valence eMonatomic
Non metallic character
decreases down the
group
Periodic Table Families Quiz
1.
Match the group number with the group name:
A. Group 1
1. Alkaline Earth Metals
B. Group 2
2. Alkali Metals
C. Group 17
3. Noble Gases
D. Group 18
4. Transition Metals
E. Groups 3-12
5. Halogens
Periodic Table Families Quiz
2. Name the group with the following properties:
a. The most reactive metals
b. The most reactive non metals
c. Monatomic gases
d. Create colored compounds and solutions
e. Have two valence and a +2 charge.
Octet Rule
 Atoms will gain or lose electrons in
order to form a stable noble gas
configuration with 8 electrons in
the valence shell
 Hydrogen, Lithium, Beryllium and
boron have 2 valence electrons like
helium
 Metals will lose valence, nonmetals
will gain.
Atomic Radius
 The size of the atom in the ground state.
Across a period, the radius decreases because more
protons pull on electrons.
 Down a group, the radius increases due to more energy
levels.

 Example: Predict the trend in radius for Be, Mg, Ca
and Sr. Then predict the trend for N, O, F and Ne.
Atomic radius trends
Ionic Radius
increases
 When atoms gain electrons, the size ___________.
decreases
 When the atom loses electrons, the size ________.
 Explain why these rules make sense.
Review of Bohr diagrams:
 Draw the Bohr diagrams of the
following atoms:
 Sodium
 Argon
 Potassium
 Fluorine
Ionization Energy
 The energy needed to remove an electron from the
valence shell of a gaseous atom or ion in the ground
state.
Across a period, IE increases because electrons are more
tightly bound to the nucleus (more p+ in the nucleus).
 Down a group, IE decreases because electrons are farther
from the nucleus and not held as tight.
 Values are found on Table S.

Ionization Energy
Table S
 Which element has the highest IE?
 Which element has the lowest IE?
 What do you think the second and third IE means?
Electronegativity
 Ability of an atom to attract electrons.
Across a period, electronegativity increases due to
stronger nucleus (number of protons).
 Down a group, electronegativity decreases due to more
principle energy levels, electrons are far from the nucleus
and can’t be attracted as well.

 These values range from 0-4 and are not energy
values.
 Notice the trend and explanations are similar to IE.
Electronegativity
Table S
 Which element has the highest e-negativity?
 Which element has the lowest e-negativity?
 Do noble gases have e-negativity values? Why?
Review
1.
Which is larger?
a. Na or Li
b. Sr or Mg
c. Ce or Ge
d. O or S
e.
f.
g.
e.
Na
I
S
Al
or
or
or
or
Na+
IS-2
Al+3
2.
Which has a higher e-negativity?
a. Cl or F
c. Mg or Ne
b. C or N
d. As or Ca
3.
Which has a greater attraction for electrons?
a. Ca or O
b. O or F
More Review
1.
Which has a higher IE?
a. Li or B
b. Mg or Sr
2. Why is the second IE always higher than the first?
3. What happens to the radius of an atom that gains
an electron? What are these ions called?