The Periodic Table and Periodic
Law
Chapter 6
History of the Periodic Table’s Development
• Late 1790s: Lavoisier compiled a list
of the 23 known elements
• 1870: 70 known elements
• 1864: John Newlands proposed an
organizational scheme for the
elements
– When elements are arranged by
increasing atomic mass their
properties repeat every eight
element.
– Named this periodic property the
law of octaves.
History of the Periodic Table’s Development
•
•
1869: Meyer and Mendeleev each demonstrated a
connection between atomic mass and elemental
properties.
– Mendeleev published his scheme first and thus
received more credit.
– Mendeleev arranged elements in columns of
increasing atomic mass since the elements had
similar properties… this developed into the first
periodic table.
– Accepted because he predicted the existence
and properties of several undiscovered
elements and left blank spaces for these
(Scandium, Gallium, and Germanium).
1913: Moseley discovered that atoms of each
element contain a unique number of protons in
their nuclei (atomic number).
– Arranging elements by increasing atomic
number correct any errors in Mendeleev’s
periodic table.
– Resulted in a clear periodic pattern of
properties (periodic law).
The Modern Periodic Table
• Boxes contain elemental name, symbol,
atomic number, and atomic mass.
• Arranged in columns (groups) and rows
(periods) by increasing atomic number.
• Groups 1A-8A (groups 1-2 and 13-18) are the
main group elements or representative
elements. S and p block
• Groups 1B-8B (groups 3-12) are the transition
elements.
The Modern Periodic Table
• Three main types of elements:
– Metals: shiny, solids at room temperature, good
conductors, ductile, malleable.
• Alkali Metals (group 1), Alkaline Earth Metals (group 2),
transition metals (groups 3-12), and inner transition metals.
– Nonmetals: gases or brittle-dull solids, poor
conductors.
• Halogens (group 17) and Nobles Gases (group 18).
– Metalloids (Semimetals): properties of metals and
nonmetals.
Classification of the Elements
• Organizing the Elements by Electron
Configuration
– Atoms in the same group have similar chemical
properties because they have the same number of
valence electrons.
• The energy level of an element’s valence electrons
indicates the period on the periodic table in which it is
found.
• A representative element’s group number and the
number of valence electrons it contains are also related
Classification of the Elements
• The s-,p-,d-, and f-Block Elements
– s-block elements: groups 1 (1 valence electron) and
group 2 (2 valence electrons).
– p-block elements: groups 13-18 (valence electrons
from 3-8).
– d-block elements: transition metals (groups 3-12).
– f-block elements:
inner transition metals.
Periodic Trends
• Many properties of elements tend to change
in a predictable way, known as a trend, as you
move across a period or down a group.
– Atomic Radius
– Ionic Radius
– Ionization Energy
– Electronegativity
Atomic Radius
• Atomic size is defined by how closely an atom lies to a
neighboring atom.
• Atomic radius is defined as half the distance between nuclei
of identical atoms that are chemically bonded together.
– Atomic Size/Radius decreases as you move left to right
across the period… this is because there is an increased
nuclear charge (more protons pulling the outer electrons
closer to the nucleus).
– Atomic Size/Radius increases as you move top to bottom
down the group… this is because more energy levels are
being added between the nucleus and the outermost
energy level (this results in less of a pull on the valence
electrons by the nucleus).
Ionic Radius
• Atoms gain or lose one or more electrons to form ions.
– An ion is an atom or group of atoms with a positive or negative charge.
• When atoms lose electrons to form a positive charge they
become smaller.
– This is because the lost electron is a valence electron… often resulting in an
empty orbital.
– There are also less electrons being pulled by the positive nucleus… which
means the electrons are pulled even closer.
• When atoms gain electrons to form a negative charge they
become larger.
– These extra electrons force the outermost electrons further from the nucleus
resulting in less of an attraction of the nucleus to the valence electrons.
• The trend for the representative elements:
– Across the period… decrease.
– Down the group… increase.
Ionization Energy
• Ionization energy is the amount of energy required
to remove an electron from an atom.
– Energy is needed to overcome the attraction between the
positive charge in the nucleus and the negative charge of
the electron.
– Think of ionization energy as an indication of how strongly
an atom’s nucleus holds onto its valence electrons.
• High ionization energy = strong hold (*nonmetals)
• Low ionization energy = weak hold (*metals)
• Second Ionization Energy: energy required to remove
a second electron.
– Third Ionization Energy… and so on
Ionization Energy
• First Ionization Energy increases as
you move left-to-right across a
period because of increased
nuclear charge.
• First Ionization Energy decreases
as you move down the group
because the valence electrons are
further from the nucleus and
easier to remove.
Electronegativity
• The ability of an atom to attract electrons in a
chemical bond.
– Fluorine is the most electronegative atom.
• Decreases as you move down the group and
increases as you move across the period.
The Periodic Table and Periodic
Law
Chapter 6
History of the Periodic Table’s Development
• Late 1790s: Lavoisier compiled a list
of the __________elements
• 1870: __________elements
• 1864: John Newlands proposed an
organizational scheme for the
elements
– When elements are arranged by
__________________their
properties repeat every eight
element.
– Named this periodic property the
____________.
History of the Periodic Table’s Development
• 1869: ________________each demonstrated a
connection between atomic mass and elemental
properties.
– ___________ published his scheme first and
thus received more credit.
– Mendeleev arranged elements in columns of
increasing atomic mass since the elements had
_______________… this developed into the
first periodictable.
– Accepted because he predicted the existence
and properties of several ________________
elements and left blank spaces for these
(Scandium, Gallium, and Germanium).
• 1913: Moseley discovered that atoms of each
element contain a unique number of protons in
their nuclei (_______________).
– Arranging elements by increasing atomic
number correct any errors in Mendeleev’s
periodic table.
– Resulted in a clear periodic pattern of
properties (_______________).
The Modern Periodic Table
• Boxes contain elemental name, symbol, atomic
number, and atomic mass.
• Arranged in columns (__________) and rows
(_________) by increasing atomic number.
• Groups 1A-8A (groups 1-2 and 13-18) are the
main group elements or _____________
elements.
• Groups 1B-8B (groups 3-12) are the ________
elements.
The Modern Periodic Table
• Three main types of elements:
– _________: shiny, solids at room temperature,
good conductors, ductile, malleable.
• Alkali Metals (group 1), Alkaline Earth Metals (group 2),
transition metals (groups 3-12), and inner transition
metals.
– _________: gases or brittle-dull solids, poor
conductors.
• Examples: Halogens (group 17) and Nobles Gases
(group 18).
– __________ (Semimetals): properties of metals
and nonmetals.
Classification of the Elements
• Organizing the Elements by Electron
Configuration
– Atoms in the same group have similar chemical
properties because they have the same number of
________________.
• The energy level of an element’s valence electrons
indicates the _______on the periodic table in which it is
found.
• A representative element’s ________ number and the
number of _____ electrons it contains are also related.
Classification of the Elements
• The s-,p-,d-, and f-Block Elements
– _________ elements: groups 1 (1 valence electron)
and group 2 (2 valence electrons).
– _________ elements: groups 13-18 (valence
electrons from 3-8).
– _________ elements: transition metals (groups 3-12).
– _________ elements:
inner transition metals.
Periodic Trends
• Many properties of elements tend change in a
predictable way, known as a trend, as you
move across a period or down a group.
– _________________
– _________________
– _________________
– _________________
Atomic Radius
• ____________is defined by how closely an atom lies to a
neighboring atom.
• ____________is defined as half the distance between nuclei
of identical atoms that are chemically bonded together.
– Atomic Size/Radius _________ as you move left to right
across the period… this is because there is an increased
nuclear charge (more protons pulling the outer electrons
closer to the nucleus).
– Atomic Size/Radius ________ as you move top to bottom
down the group… this is because more energy levels are
being added between the nucleus and the outermost
energy level (this results in less of a pull on the valence
electrons by the nucleus).
Ionic Radius
• Atoms gain or lose one or more electrons to form _____.
– An ion is an atom or group of atoms with a positive or negative charge.
• When atoms lose electrons to form a positive charge they
become _________.
– This is because the lost electron is a valence electron… often resulting in
an empty orbital.
– There are also less electrons being pulled by the positive nucleus…
which means the electrons are pulled even closer.
• When atoms gain electrons to form a negative charge they
become _________.
– These extra electrons force the outermost electrons further from the
nucleus resulting in less of an attraction of the nucleus to the valence
electrons.
• The trend for the representative elements:
– ______________________… decrease.
– ______________________… increase.
Ionization Energy
• _____________________is the amount of energy
required to remove an electron from an atom.
– Energy is needed to overcome the attraction between the
positive charge in the nucleus and the negative charge of
the electron.
– Think of ionization energy as an indication of how strongly
an atom’s nucleus holds onto its valence electrons.
• High ionization energy = strong hold
• Low ionization energy = weak hold
• _________________________: energy required to
remove a second electron.
– Third Ionization Energy… and so on
Ionization Energy
• First Ionization Energy __________
as you move left-to-right across a
period because of increased
nuclear charge.
• First Ionization Energy __________
as you move down the group
because the valence electrons are
further from the nucleus and
easier to remove.
Electronegativity
• The ability of an atom to attract electrons in a
chemical bond.
– ________ is the most electronegative atom.
• ____________ as you move down the group
and _______ as you move across the period.