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Exploring the Periodic Table
Modern Chemistry; Holt, Rinehart, & Winston
CHAPTER 5 – SECTION 1
HISTORY OF THE PERIODIC TABLE
In the late 1800s, scientists had identified over 60 elements.
Certain characteristic physical and chemical properties were
associated with each element. The physical property called atomic
mass provided chemists with a convenient way to organize the
elements. At the same time, it was recognized that there were
certain elements that had similar chemical properties. Mendeleev
arranged the elements in rows according to atomic weight and
kept elements with similar chemical properties in the same
columns. Today elements are ordered according to atomic number
rather than atomic mass.
Learning Targets




I can explain the roles of Mendeleev and
Moseley in the development of the periodic
table.
I can describe the modern periodic table.
I can explain how the periodic law can be
used to predict the physical and chemical
properties of elements.
I can describe how the elements belonging
to a group of the periodic. table are
interrelated in terms of atomic number.
Stanislao Cannizzaro (1826-1910)



Italian chemist
Determined a method
for accurately measuring
the relative masses of
atoms
His method allowed
chemists to search for a
relationship between
atomic mass and other
properties of elements
Dmitri Mendeleev (1834-1907)




Russian chemist
Credited as being the creator of the
first version of the periodic table of
elements
Arranged his periodic table
according to atomic mass so that
elements with similar properties
were in the same group

Some elements could not be
arranged according to atomic
mass in order to keep the
elements arranged according
to properties
Predicted the properties of
elements that had not yet been
discovered using his periodic table
Mendeleev’s Periodic Table
“I began to look about and write down the elements with their atomic weights and typical
properties, analogous elements and like atomic weights on separate cards, and this soon convinced
me that the properties of elements are in periodic dependence upon their atomic weights.”
--Mendeleev, Principles of Chemistry, 1905, Vol. II
Henry Moseley (1887-1915)



English chemist
Worked with Rutherford
Proved Mendeleev’s
arrangement of the periodic
table to be correct – only, the
periodic table was arranged
according to atomic number,
not atomic mass
The Periodic Law

States that when elements are arranged
in order of increasing atomic number,
their physical and chemical properties
show a periodic pattern
CHAPTER 5 – SECTION 2
ELECTRON CONFIGURATION AND THE
PERIODIC TABLE
The modern periodic table has 112 squares, which represent a unique
element. The distinctive shape of the periodic table comes in part from
the periodic law. Elements in the same column have similar properties.
These columns are referred to as groups or families of elements. The
horizontal rows of the periodic table are called periods. The elements
in the periodic table are also grouped as metals, nonmetals, and
semimetals. Metals make up most of the periodic table and are located
in the center and at the left of the table. With the exception of
hydrogen, nonmetals are on the right side, and semimetals are located
between the metals and nonmetals. The periodic table can also be
viewed in terms of orbital blocks. These orbital blocks refer to the
orbitals (s, p, d, and f ) which contain the elements’ incompleted
sublevels of electrons.
Learning Targets




I can describe the relationship between electrons in
sublevels and the length of each period of the
periodic table
I can locate and name the four blocks of the periodic
table and explain the reasons for these names
I can discuss the relationship between group
configurations and group numbers
I can describe the locations in the periodic table and
the general properties of the alkali metals, the
alkaline-earth metals, the halogens, the transition
metals, the noble gases, the actinides, the
lanthanides, the metals, the nonmetals, the
metalloids, and the main group elements
Periodic Law Demonstrated in Groups

Why do
elements in
groups have
similar physical
and chemical
properties?


They have the same
number of valence
electrons in their
outer energy levels.
Generally, the
configurations of the
outermost electron
shells of elements
within the same
group are the same.
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
METALS
METALLOIDS
NONMETALS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
ALKALI
METALS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
ALKALINE-EARTH
METALS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
HALOGENS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
NOBLE GASES
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
TRANSITION
METALS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
INNER TRANSITION
(Rare Earth)
METALS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
LANTHANIDES
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
ACTINIDES
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
PERIODS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
GROUPS
In the periodic table below, indicate the location of the
groups, periods, alkali metals, alkaline earth metals,
halogens, noble gases, lanthanides, actinides, transition
metals, inner transition metals, main group elements,
metals, nonmetals and metalloids.
MAIN GROUP
ELEMENTS
Let’s Compare!
Metals
 Good conductors
of heat and
electricity
 Malleable
 Ductile
 Luster
 Typically solids
at room
temperature
Nonmetals
 Solids, liquids
and gases at
room
temperature
 Solids are brittle
and dull
 Poor conductors
of heat and
electricity
Metalloids
 Have properties
of both metals
and nonmetals
 Mostly brittle
solids
 Intermediate
conductors of
electricity- AKA
semiconductors
Properties of Alkali Metals





Extremely reactive
 Readily react with water
and air
Silvery in appearance
Soft enough to cut with a
knife
Lower densities than other
metals
Lower melting points than
other metals
Properties of Alkaline-Earth Metals
Harder & stronger
than alkali metals
 Higher densities &
melting points than
alkali metals
 Less reactive than
alkali metals

Properties of Halogens
Most reactive
nonmetals
 React readily with
most metals to
form salts
 Most
electronegative
elements

Properties of Noble Gases

Least reactive
elements because
their highest
occupied energy
levels are completely
filled with an octet
of electrons (except
He, which only
requires 2 electrons
to be filled).
Properties of Transition Metals
High densities
 High melting points
 Good conductors of heat &
electricity
 High luster
 Less reactive than alkali
and alkaline-earth metals

Properties of p Block Metals
Harder and more dense than the s
block metals
 Softer and less dense than the d block
metals.

Properties of Lanthanides
Soft, silvery metals
 Similar reactivity to alkaline-earth
metals

Properties of Actinides
All radioactive
 The first 4 have been found naturally
on Earth

Did you know?
Oxygen, carbon, hydrogen and
nitrogen make up 96% of the human
body mass
 Calcium and phosphorous make up 3%
 Sodium, potassium, chloride and
magnesium make up 0.7%
 Iron, cobalt, copper, zinc, selenium,
cyanide and fluorine are found in trace
amounts

CHAPTER 5 – SECTION 3
ELECTRON CONFIGURATIONS AND
PERIODIC PROPERTIES
Many of the properties of the elements change in predictable ways as
you move across a period or move down a group of the periodic table.
The predictable changes in these properties are called periodic trends.
There are periodic trends for properties such as atomic radius, ionic
size, ionization energy, electron affinity, and electronegativity.
Knowledge of these trends helps develop a better understanding of
the periodic table and of the patterns of behavior of the elements.
Learning Targets




I can define the term periodic trend.
I can define atomic radius, ionic radius, ionization
energy, electron affinity and electronegativity.
I can describe the general trends on the periodic
table for atomic radius, ionic radius, electron affinity,
ionization energy and electronegativity.
I can apply the trends on the periodic table to
answer questions regarding size, electron affinity,
ionization energy and electronegativity.
Atomic Radii

Atomic radius – one-half the distance
between the nuclei of identical atoms
that are bonded together
Atomic Radius
Distance between nuclei
Period Trends

Decreases across a period
Why?
Protons are added to the nucleus
moving across a period from left to
right
 This increases the charge of the
nucleus (effective nuclear charge – Zeff)
 As Zeff increases, the electrons are
pulled closer to the nucleus

Period Trends
++
+++++
Group Trends

Increase down a group
Why?
The addition of shells increases the
electrons’ distance from the nucleus
and the size of the atom
 Electron-electron repulsion “plumps” up
n=3
the atom
n=2
 Zeff decreases the further the electrons
n=1
are from the nucleus

Variations in Atomic Radii
Atomic Radii Trends
DECREASES
DECREASES
Ionization Energy

The energy required to remove one
electron from a neutral atom of an
element creating an ion
A + Energy  A+ + e-
Period Trends
Increase across a period
 Why?


Zeff increases across the period
Group Trends
Decrease down the group
 Why?

Electron shielding causes a decrease in
effective nuclear charge
 Electron-electron repulsion forces
increase

Variations in Ionization
Energies
Draw the orbital notation for Group 5A and Group 6A.
Can you explain the dips in the chart for these 2 groups?
Variations in Ionization
Energies
If removing an electron will create an empty or ½ filled subshell,
ionization energy will decrease.
Successive Ionization Energies
Each successive electron removed from
an ion feels an increasingly stronger
effective nuclear charge (Zeff) –
therefore, successive ionization
energies are larger than 1st ionization
energies
 A large jump in ionization energy
occurs when removing an electron from
an ion that assumes a noble gas
configuration

Ionization Energy Trends
INCREASES
INCREASES
Electron Affinity

The change in energy that a neutral
atom undergoes when an electron is
acquired (the ability to attract an e -)
A + e-  A- + energy
[negative energy value (exothermic)]
A + e- + energy  A[positive energy value (endothermic)]
Period Trends
Increase across a period
 Why?


Zeff increases across the period
Group Trends
Decrease down the group
 Why?

Electron shielding causes a decrease in
effective nuclear charge
 Electron-electron repulsion forces
increase

Variations in Electron
Affinities
Electron Affinity Trends
INCREASES
INCREASES
Ionic Radii

Cation – positively charged ion
Cations are smaller than their parent atom – why?

Removal of an electron creates an unbalanced positive charge
increasing Zeff and decreasing the radius of the ion.
Anion – negatively charged ion
Anions are bigger than their parent atom – why?
Addition of an electron creates an unbalanced negative charge
decreasing Zeff and increasing the radius of the ion.
+
++ +
+
+ +
Ionic Radii Trends
DECREASES
DECREASES
Valence Electrons
Electrons available to be gained, lost or
shared in the formation of a chemical
compound
 Located in the outer energy level

Electronegativity


A measure of the ability of an atom in a
chemical compound to attract a
bonding pair of electrons
NOTE *Electronegativity is a property
of atoms in compounds and thus differs
from ionization energy and electron
affinity, which are properties of isolated
atoms*
Trends

Increase across a period
Effective nuclear charge increases

Decrease down a group
Increase in atomic size and increase in electron shielding decreases the
effective nuclear charge

Electronegativity depends upon:
 The
number of protons in the nucleus
 The distance from the nucleus
 Electron shielding
Electronegativity Trends
INCREASES
INCREASES