Chapter 6
The Periodic Table
Objectives:
 Describe the periodic tables of Moseley and
Mendeleev.
 Identify the various families of elements on the
periodic table.
 State the trends in atomic radius, ionization
energy, electron affinity and ion size with a group
or period on the periodic table.
 Identify the relationship between these trends and
the structure of the atom.
Chapter 6
Section 1
History of the
Periodic Table
By 1860, more than 60 elements had
been discovered.
Chemists needed to learn the properties
of these elements as well as those of the
compounds they formed.
At the time there was no method for
accurately determining an element’s
atomic mass.
Things were disorganized and confusing
among chemists.
Dmitri Mendeleev
Mendeleev hoped to organize the elements
according to their properties.
He arranged the elements in order of
increasing atomic mass and noticed certain
similarities in their chemical properties would
appear at regular intervals.
In 1869 Mendeleev created the first periodic
table grouping the elements with similar
properties.
Henry Moseley
In 1911 Moseley revised the periodic table
of Mendeleev by arranging the elements in
order of increasing atomic number (number
of protons) and not atomic mass.
The elements were still grouped by similar
physical and chemical properties.
Modern Periodic Table
The periodic table has undergone
extensive changes since Mendeleev’s
time.
More than 40 new elements have been
discovered or synthesized.
Periodic Table – an arrangement of the
elements in order of their atomic numbers
so that elements with similar properties fall
in the same column or group.
Periodic Table

Elements are classified in three major
groups:




Metals
Nonmetals
Metalloids
Use the periodic table to distinguish the
classes of elements
Physical Properties of Metals






Malleability
Ductility
Luster
Heat conductors and
electrical conductors
Solids
Ex. Iron (Fe), tin (Sn),
zinc (Zn), and copper
(Cu)
Properties of Nonmetals
Dull in appearance
 Brittle
 Do not conduct electricity
 Ex. Carbon (C), Oxygen (O)
and Sulfur (S)
 Solids, liquids or gases

The metalloids divide the metals from the
nonmetals.
They are mostly brittle solids with some
properties of metals and some of nonmetals.
The electrical conductivity falls between the
metals and nonmetals.
Metalloids

Properties of metals and nonmetals

Ex. Silicon (Si) and Germanium (Ge)

Common in the computer
industry
Summary
Recognize the work of Mendeleev and
Moseley.
How the modern periodic table is
arranged with respect to the elements.
Know how the periodic table is divided
into different groups.
Homework
Page 198 (worksheet)
Questions 29, 30 , 31 and 32
Chapter 6
Section 2
Electron Configuration
and the
Periodic Table
The elements arranged vertically in the
periodic table share chemical properties due
to the same number of valence electrons.
They are also organized in horizontal rows,
or periods.
The length of each period is determined by
the number of electrons that can occupy the
sublevels in that period.
The periodic table is divided into 4 blocks
The s-Block Elements
The elements of the s block are chemically
reactive metals.
The elements of the s-block and p-block are
known as the main-group elements.
The outermost energy level in an atom of the
Group 1 elements contain a single electron.
The ease with which the single electron is lost
make the Group 1 metals extremely reactive.
Alkali Metals – The elements in Group 1 of
the periodic table (lithium, sodium, potassium,
rubidium, cesium and francium).
Because the alkali metals are so reactive they
are not found in nature as free elements.
They combine vigorously with most
nonmetals.
They react strongly with water to produce
hydrogen gas.
Alkaline-Earth Metals – The elements in
Group 2 of the periodic table (beryllium,
magnesium, calcium, strontium, barium and
radium).
Atoms of the alkaline-earth metals contain a
pair of electrons in their outermost shell.
Slightly less reactive than the alkali metals.
Still too reactive to be found in nature as free
elements.
The d-Block Elements
The elements of the d block are known as the
transition metals.
They are good conductors of electricity and
have a high luster.
They are typically less reactive than the alkali
and alkaline-earth metals.
Palladium, platinum and gold are among the
least reactive of all elements.
The p-Block Elements
The elements of the p-block and s-block are
known as the main-group elements.
The p-block consists of the elements in Groups
13-18 (except helium).
The properties of the p-block vary greatly.
The p-block contains metals (Al), metalloids
(Si) and nonmetals (Br).
Halogens – the elements of Group 17
(fluorine, chlorine, bromine, iodine and
astatine).
The halogens are the most reactive
nonmetals.
They react vigorously with most metals to
form compounds known as salts (NaCl).
Fluorine and chlorine are gases. Bromine is
a liquid and iodine is a solid.
Noble Gases (1868-1900) – group 18
elements that are characterized by their
relative un-reactivity.
The Lanthanides (early 1900’s) – the 14
elements with atomic numbers from 58
(cerium, Ce) to 71 (lutetium, Lu)
The Actinides – the 14 elements with
atomic numbers from 90 (thorium, Th) to
103 (lawrencium, Lr).
The lanthanides and actinides belong in
periods 6 and 7 of the periodic table.
To save space and to group them
together they are set off below the main
portion of the periodic table.
The Group 18 elements (Noble Gases)
undergo few chemical reactions.
This stability results from the gases
electron configuration.
Their highest occupied levels are
completely filled with electrons (octet).
An atom’s electron configuration governs
the atom’s chemical properties.
Homework
Page 198 (worksheet)
Questions 43, 44, 47 and 48
Chapter 6
Section 3
Periodic Trends
So far you have learned that the elements
are arranged in the periodic table according
to their atomic number.
There is also a correlation between the
arrangement of the elements and their
electronic configuration.
We will look at the relationship between the
electron configurations and the periodic
trends of the elements.
Atomic Radii
Ideally, the size of an atom is defined by
the edge of its orbital.
However, this boundary varies under
different conditions.
One way to express the atomic radius is to
measure the distance between the nuclei
of two identical atoms that are bonded
together, then divide this distance by two.
Atomic radius may be defined as one-half the
distance between the nuclei of identical atoms
that are bonded together.
There is a gradual decrease across a row.
The trend to smaller atoms across a period is
caused by increasing positive charge of the
nucleus. Adding of protons or increasing
atomic number.
The electrons are pulled closer to the
nucleus.
The increase pull results in a smaller atomic
radius.
Group Trends
There is an increase down a group.
There is an increase down a group.
As electrons are added to sublevels in
higher energy levels located further from the
nucleus, the size of the atoms increase.
An exception is between aluminum (radius –
143 pm) and gallium (radius – 135 pm).
This is due to the gallium being proceeded
by the 10 d-block elements.
The nuclear charge is considerably higher.
Problem
Of the elements magnesium-Mg, chlorineCl, sodium-Na, and phosphorus-P, which
has the largest atomic radius and why?
Of the elements calcium-Ca, beryllium-Be,
barium-Ba and strontium-Sr, which as the
largest atomic radius and why?
Classwork
Page 189
Problems: 16 and 17
Ionization Energy
Ionization energy (IE) – the energy
required to remove one electron from a
neutral atom of an element.
A + energy
A+ + e-
An ion is an atom or group of bonded
atoms that has a positive or negative
charge.
Period Trends
In general, ionization energies of the
elements increase across a period.
Group 1 elements – have the lowest
ionization energies. Therefore they lose
electrons most easily. Very reactive.
Group 18 elements - have the highest
ionization energies . They do not lose
electrons easily. Very low reactivity.
The increase is due to increasing the number
of protons going across a period.
Increasing the number of protons more
strongly attracts electrons in the same energy
level.
Therefore, it is tougher to remove an electron
from an atom.
Increasing proton number is responsible for
both an increasing ionization energy and
decreasing atomic radius across a period.
Group Trends
Ionization energies generally decrease down
a group.
Electrons removed from atoms of the
elements down a group are farther from the
nucleus.
Also, the electrons from the lower energy
levels shield the outer electrons.
Therefore, they are removed more easily.
Ions
Atoms can gain or lose electrons to form ions.
An ion is an atom that has either a positive or
negative charge.
A
A + e-
A+ + e-
(loss of electron)
A-
(gain of electron)
Cations and Anions
Cation (A+) – a positive ion that results with
loss of an electron.
Anion (A-) – a negative ion that results with
the gain of an electron.
Electronegativity
Electronegativity is a measure of the ability
of an atom in a chemical compound to
attract electrons.
Period Trends
Electronegativities tend to increase across a
period.
The Group 1 and 2 elements have the lowest
electronegativities.
Groups 16 and 17 have the highest
electronegativities.
Group Trends
Electronegativities tend to decrease down a
group.
The combination of these trends in
electronegativities results in the highest
values belonging to the elements in the
upper right of the periodic table (fluorine).
Problem
Of the elements gallium-Ga, bromine-Br,
and calcium-Ca, which has the highest
electronegativities and why?
Valence Electrons
Chemical compounds form because
electrons are lost, gained or shared
between atoms.
The electrons that interact are those in the
highest energy levels.
Valence electrons – the electrons available
to be lost, gained or shared in the
formation of chemical compounds.
Valence electrons are located in s and p
blocks of the periodic table.
The valence electrons are located on the
same row as the element.
For example, the electron lost in Na to form
Na+ is from the 3s sublevel.
Ignore the d block elements.
The inner electrons are in filled energy
levels and are held to tightly to be involved.
Homework
Chapter Review
Page 199
Questions 58, 60, 63 and 64