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Unit 3: Periodic Table Chapter 6
Objectives
21 Understand the historical background of the periodic table including
such contributions of Newlands, Mendeleev, and Moseley
22 Use the periodic table to predict properties of certain elements
23 Identify the difference between periods and groups
24 Identify where main group, metals, metalloids, non-metals, alkali,
alkaline-earth, halogen, noble gases, transition metals, lanthanides
(rare earth metals) and actinides exist on the periodic table along
with their characteristics and electron configurations
25 Define and apply the periodic law to the trends on the periodic table
including atomic radius, ionization energy, and electron
affinity/electronegativity
21 Historical Background on the
Periodic Table The Periodic Table was original designed by
John Newlands and was published in the
1860s.
◦ He also proposed an idea known as the Law of
Octaves to help explain his setup for the table.
More extensive work was done on the
Periodic Table by Dmitri Mendeleev.
◦ He arranged the table by atomic mass.
◦ He left blanks where he felt elements should
belong even though they were not discovered
yet.
Mendeleev’s Periodic Table
Historical Background Continued
While Mendeleev’s table was the first that
resembled our current model, there were
some flaws.
◦ None of the noble gases were present.
◦ There were some discrepancies Mendeleev could
not explain in certain properties.
◦ It did not account for isotopes.
Historical Background Continued
In 1914, Henry Moseley rearranged the
Periodic Table by atomic number.
By doing this, he eliminated the
discrepancies that Mendeleev could not
explain.
Moseley’s table is the model used today with
the elements found after 1914 added to the
table.
22 Characteristics of the Table
The Periodic Table was carefully designed to
provide as much information as possible.
The table is ordered into:
◦ Periods: horizontal rows
◦ Groups: vertical columns
◦ Blocks: S, P, D, or F
23,24 Sections of the Table
There are certain sections of the Periodic
Table that have common properties.
◦ Metals
◦ Metalloids
◦ Non-metals
◦ Main Group Elements
Metals
Shiny
Form positive ions
Ductile
Malleable
High melting and
boiling points
Good conductors
◦ Heat and energy
Return
Non-Metals Make up the majority
of the crust, atmosphere and living organisms
Low melting and boiling points
Form negative ions
Low densities
If solid, tend to be dull and brittle
Poor conductors
◦ Both heat and energy Return
Metalloids
Semi-conductors
Contain some
metallic properties
Contain some non-
metallic properties
Return
Main Group Elements
Made up of the S and
P blocks
Consists of some of
the most common
elements.
Return
Groups and Blocks
Alkali Metals
Alkaline Earth Metals
Transition Metals
Halogens Noble Gases
Actinide Series
Lathanide Series
Alkali Metals
Highly reactive
Rarely found in the
elemental form
Soft metals
Low densities
Make +1 ions
Last electron is
always a s1
Return
Alkaline Earth Metals
Always from +2 ions
Last electron is always an s2
High melting points
Reactive but not as violent as the alkali metals
Return
Halogens
Highly reactive
Form -1 ions
All but astatine can
form a diatomic
molecule
Common in acids
Used as disinfectants
and in pesticides
Return
Noble Gases
Odorless, colorless
gases
Outer (valence)
energy level is full
Very low reactivity
Melting and boiling
points are low and
very close together
Cryogenic
refrigerants Return
Transition Metals
Form the D-Block of
the Periodic Table
Magnetic Properties
◦ 1 or more unpaired
electrons
High melting and boiling
points
Can from +1, +2, +3
ions
Generally solid
Return
Lanthanide Series
Make up the 4f block
Typically used in lasers
Sometimes referred to
as the rare earth
metals
◦ Though actually found in
high concentrations in
the crust
Superconductors
Batteries and magnets
Return
Actinide Series
Make up the 5f-block
Most are man-made
◦ Thorium and uranium are
the only two the occur
naturally with any
abundance.
Radioactive
Return
25 Periodic Law
Certain properties follow periodic law.
Periodic law refers to the increasing or
decreasing of a trend as one progresses
across a period or group on the Periodic
Table.
Three of the most common trends that are
monitored and follow periodic law are
atomic radius, ionization energy, and
electron affinity.
Atomic Radius
Atomic radius refers to the size of the
electron cloud surrounding the nucleus.
The atomic radius increases as each new
energy level is added.
While electrons are being added to an
energy level, electron shielding allows for the
affects the size of the atom.
Electron Shielding
As electrons start filling energy levels, the nucleus holds them close.
As more are added, the inner ring prevents the nucleus from pulling the outer ring too close (it shields the positive charge).
The nucleus will pull the energy level slightly closer though as you progress across the table.
Atomic Radius Trend
The atomic radius increases in the
direction of the arrow.
Ionization Energy
Ionization energy is the energy required to
remove an electron from an atom.
The larger the atom, the more difficult it is
for the nucleus to hold onto its electrons.
Smaller atoms can hold onto electrons much
easier.
Ionization Energy Trend
The ionization energy increases in the
direction of the arrow.
Electronegativity
Electronegativity refers to how well an
atom attracts electrons.
Smaller atoms have more nuclear charge
to attract electrons.
As that large atoms have a difficult time
holding onto their electrons, they do not
readily attract electrons.
Electronegativity Trend
The electronegativity increases in the
direction of the arrow.
This concludes the tutorial on
measurements.
To try some practice problems, click here.
To return to the objective page, click
here.
To exit the tutorial, hit escape.
Definitions-Select the word to return to the tutorial