Chapter 6
The Periodic Table and
Periodic Law
The elements, which make up all
living and non-living matter, fit into a
orderly table. When interpreted
properly, the table describes much of
the elements physical and chemical
properties.
What is the Periodic Law and
how was it formulated?
• Demitri Mendeleev is
known as the father of
the periodic table
• He arranged the
elements in families
(groups) and periods
(rows,series)
according to atomic
mass and properties
• Mendeleev noted that the chemical and
physical values for elemental properties
would either be high or low depending upon
the group under observation.
• He proposed the first Periodic Law "The
properties of the elements are a periodic
function of their atomic masses"
• Left blanks in his table for undiscovered
elements
Moseley’s Modern-Periodic Law
• There was some
inconsistencies with
Mendeleev’s table
• In the early 1900’s
Moseley was able to
experimentally
determine the atomic
number of all known
elements
• Moseley then proceeded to rearrange the
elements according to increasing atomic
numbers.
• New/Modern Periodic law states that the
properties of elements are a periodic
function of their atomic number
The Modern Periodic Table
• Glenn T. Seaborg won
the Nobel Prize for his
work in nuclear
chemistry
• In 1944, formulated
the “actinide concept”
of heavy element
electronic structure.
This concept predicted
that the fourteen
actinides
Some Characteristics of Groups
Group 1 (IA) - Alkali Metals
( metal characteristics - shiny, malleable,
ductile, good conductors)
• Very active metals - activity increases as
you go down group
• All have one valence electron - form +1
cations by losing an electron
• react violently with water
Group 2 (IIA) - Alkaline Earth
Metals
• activity increases as you move down
the column not as reactive as alkali
metals
• Ca, Sr, and Ba react violently when
they come into contact with water
• All have two valence electrons
• Form +2 cations by losing 2 electrons
Group 17 (VIIA) - Halogens
• All gain one electron to form anions with a
charge of -1.
• All are nonmetals except for At which is a
semimetal
• All are diatomic in their elemental form
Group 18 (VIIIA) - Noble (Rare)
Gases
• Mistakenly labeled as "inert gases" until
about 30 years ago because it was thought
that these gases did not react with anything.
• Noble gases have filled valence (outermost)
shells.
Periodic Trends
• Atomic Radii
1) As you move down a group, atomic
radius increases.
2) WHY? - The number of energy levels
increases as you move down a group as the
number of electrons increases.
• As you move across a period, atomic
radius decreases.
• WHY? - As you go across a period,
electrons are added to the same energy
level. At the same time, protons are being
added to the nucleus. The concentration of
more protons in the nucleus creates a
"higher effective nuclear charge."
First Ionization Energy
• Definition: The energy required to remove
the outermost (highest energy) electron
from a neutral atom in its ground state.
1) As you move down a group, first
ionization energy decreases.
2) WHY? Electrons are further from the
nucleus and thus easier to remove the
outermost one - + shielding effect of
other electrons
3) As you move across a period, first
ionization energy increases.
4) WHY? - As you move across a period,
the atomic radius decreases, the
attraction from the positive nucleus gets
larger
Electronegativity
• Definition: The ability of an element to
attract electrons in a chemical bond
1) As you move down a group,
electronegativaty decreases
2) As you move across a period, it
increases.