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Transcript
The Periodic Table and Periodic Law
Section 6.1 Development of the
Modern Periodic Table
Section 6.2 Classification of the
Elements
Section 6.3 Periodic Trends
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Section 6.1 Development of the Modern
Periodic Table
• Trace the development of the periodic table.
• Identify key features of the periodic table.
atomic number: the number of protons in an atom
The periodic table evolved over
time as scientists discovered more
useful ways to compare and
organize the elements.
Section 6.1 Development of the Modern
Periodic Table (cont.)
periodic law
group
period
representative elements
transition elements
metal
alkali metals
alkaline earth metals
transition metal
inner transition metal
lanthanide series
actinide series
nonmetals
halogen
noble gas
metalloid
Development of the Periodic Table
• In the 1700s, Lavoisier compiled a list of all
the known elements of the time.
Development of the Periodic Table (cont.)
• The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
• John Newlands proposed an arrangement
where elements were ordered by increasing
atomic mass.
Development of the Periodic Table (cont.)
• Newlands noticed
when the elements
were arranged by
increasing atomic
mass, their
properties repeated
every eighth
element.
Development of the Periodic Table (cont.)
• Meyer and Mendeleev both demonstrated
a connection between atomic mass and
elemental properties.
• Moseley rearranged the table by increasing
atomic number, and resulted in a clear
periodic pattern.
• Periodic repetition of chemical and physical
properties of the elements when they are
arranged by increasing atomic number is
called periodic law.
Development of the Periodic Table (cont.)
The Modern Periodic Table
• The modern periodic table contains boxes
which contain the element's name, symbol,
atomic number, and atomic mass.
The Modern Periodic Table (cont.)
• Columns of elements are called groups.
• Rows of elements are called periods.
• Elements in groups 1,2, and 13-18 possess a
wide variety of chemical and physical
properties and are called the representative
elements.
• Elements in groups 3-12 are known as the
transition metals.
Periods
• Each row is a period
• The number of elements per period varies due to the
way the orbitals increase from NRG level to NRG level
Groups
• Each column in the periodic
table is called a group
• Elements within a group
have similar properties
• Elements in a group have
similar electron
configurations
• The electron configuration
of an element determines
its chemical properties
Classes of Elements
• 3 ways to classify elements
– States at room temperature: solid, liquid, gas
– Naturally occurring or not naturally occurring
(atomic #93 and higher do not occur naturally)
– Categories based on their general properties
• Metals – located on the left in the periodic table
• Non metals – located on the right
• Metalloids – in-between
Metals
• Most on the periodic table are classified as metals
• Properties
– Good conductors
– Except mercury, they are solid at room temp
– Malleable
– Ductile
Groups 3-12 are transition metals
can form compounds with distinctive colors
Non Metals
• Properties on
nonmetals
– Opposite of metals
– Poor conductors
– Many are gases at
room temp
– The nonmetals
that are solid at
room temp are
very brittle
Metalloids
• Properties of metalloids
– Between metals and nonmetals
– A metalloid’s ability to conduct electric current
varies depending on temperature
The Modern Periodic Table (cont.)
• Elements are classified as metals,
non-metals, and metalloids.
• Metals are elements that are generally shiny
when smooth and clean, solid at room
temperature, and good conductors of heat
and electricity.
• Alkali metals are all the elements in group 1
except hydrogen, and are very reactive.
• Alkaline earth metals are in group 2, and
are also highly reactive.
The Modern Periodic Table (cont.)
• The transition elements are divided into
transition metals and inner transition
metals.
• The two sets of inner transition metals are
called the lanthanide series and actinide
series and are located at the bottom of the
periodic table.
The Modern Periodic Table (cont.)
• Group 17 is composed of highly reactive
elements called halogens.
• Group 18 gases are extremely unreactive and
commonly called noble gases.
The Modern Periodic Table (cont.)
• Metalloids have physical and chemical
properties of both metals and non-metals,
such as silicon and germanium.
The Modern Periodic Table (cont.)
Section 6.1 Assessment
What is a row of elements on the periodic
table called?
A. octave
B. period
D
A
0%
C
D. transition
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. group
Section 6.1 Assessment
What is silicon an example of?
A. metal
B. non-metal
C. inner transition metal
D
C
A
0%
B
D. metalloid
A. A
B. B
C. C
0%
0%
0%
D. D
Section 6.2 Classification of the Elements
• Explain why elements in
the same group have
similar properties.
• Identify the four blocks
of the periodic table
based on their electron
configuration.
valence electron:
electron in an atom's
outermost orbitals;
determines the chemical
properties of an atom
Elements are organized into different
blocks in the periodic table according
to their electron configurations.
Organizing the Elements by Electron
Configuration
• Recall electrons in the highest principal
energy level are called valence electrons.
• All group 1 elements have one valence
electron.
Organizing the Elements by Electron
Configuration (cont.)
• The energy level of an element’s valence
electrons indicates the period on the
periodic table in which it is found.
• The number of valence electrons for
elements in groups 13-18 is ten less than
their group number.
Organizing the Elements by Electron
Configuration (cont.)
The s-, p-, d-, and f-Block Elements
• The shape of the periodic table becomes
clear if it is divided into blocks representing
the atom’s energy sublevel being filled with
valence electrons.
The s-, p-, d-, and f-Block Elements (cont.)
• s-block elements consist of group 1 and 2,
and the element helium.
• Group 1 elements have a partially filled s
orbital with one electron.
• Group 2 elements have a completely filled s
orbital with two electrons.
The s-, p-, d-, and f-Block Elements (cont.)
• After the s-orbital is filled, valence
electrons occupy the p-orbital.
• Groups 13-18 contain elements with
completely or partially filled p orbitals.
The s-, p-, d-, and f-Block Elements (cont.)
• The d-block contains the transition metals
and is the largest block.
• There are exceptions, but d-block elements
usually have filled outermost s orbital, and
filled or partially filled d orbital.
• The five d orbitals can hold 10 electrons, so
the d-block spans ten groups on the periodic
table.
The s-, p-, d-, and f-Block Elements (cont.)
• The f-block contains the inner transition
metals.
• f-block elements have filled or partially filled
outermost s orbitals and filled or partially filled
4f and 5f orbitals.
• The 7 f orbitals hold 14 electrons, and the
inner transition metals span 14 groups.
Section 6.2 Assessment
Which of the following is NOT one of the
elemental blocks of the periodic table?
A. s-block
B. d-block
D
A
0%
C
D. f-block
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. g-block
Section 6.2 Assessment
Which block spans 14 elemental groups?
A. s-block
B. p-block
C. f-block
D
C
A
0%
B
D. g-block
A. A
B. B
C. C
0%
0%
0%
D. D
Section 6.3 Periodic Trends
• Compare period and
group trends of several
properties.
• Relate period and group
trends in atomic radii to
electron configuration.
principal energy level:
the major energy level of
an atom
ion
ionization energy
octet rule
electronegativity
Trends among elements in the periodic
table include their size and their ability
to lose or attract electrons
• Because of the way the periodic table is
organized today, we can extrapolate a lot
of information about an element by looking
at its location on the table
– Atomic radii
– Ionization energy
– Metallic Character
– Non-metallic Character
– Electronegativity
Atomic Radii
• Measure of the size of an element’s atoms
• Distance from the nucleus to the surrounding
outer edge of the cloud of electrons
• Across a period from left to right the atomic
radii will decrease
• From the top to the bottom of a group, the
atomic radii will increase
• Draw arrows on your periodic table to
indicate these trends
Atomic Radius
• Atomic size is a periodic trend influenced
by electron configuration.
• For metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the
element.
Atomic Radius (cont.)
• For elements that occur as molecules, the
atomic radius is half the distance between
nuclei of identical atoms.
Atomic Radius (cont.)
• There is a general decrease in atomic
radius from left to right, caused by
increasing positive charge in the nucleus.
• Valence electrons are not shielded from the
increasing nuclear charge because no
additional electrons come between the
nucleus and the valence electrons.
Atomic Radius (cont.)
Atomic Radius (cont.)
• Atomic radius generally increases as you
move down a group.
• The outermost orbital size increases down a
group, making the atom larger.
Ionization Energy (cont.)
• The octet rule states that atoms tend to
gain, lose or share electrons in order to
acquire a full set of eight valence electrons.
• The octet rule is useful for predicting what
types of ions an element is likely to form.
Ionic Radius
• An ion is an atom or bonded group of
atoms with a positive or negative charge.
• When atoms lose electrons and form
positively charged ions, they always become
smaller for two reasons:
1. The loss of a valence electron can leave an empty
outer orbital resulting in a small radius.
2. Electrostatic repulsion decreases allowing the
electrons to be pulled closer to the radius.
Ionic Radius (cont.)
• When atoms gain electrons, they can
become larger, because the addition of an
electron increases electrostatic repulsion.
Ionic Radius (cont.)
• The ionic radii of positive ions generally
decrease from left to right.
• The ionic radii of negative ions generally
decrease from left to right, beginning with
group 15 or 16.
Ionic Radius (cont.)
• Both positive and negative ions increase in
size moving down a group.
Ionization Energy
• Ion = an atom that has lost or gained an
electron causing it to have a positive or
negative charge
• Ionization Energy = The energy required to
remove the outermost electron from an
atom
• If an electron is pulled away from the
atom, an ion is formed
Ionization Energy
• Ionization energy is defined as the energy
required to remove an electron from a
gaseous atom.
• The energy required to remove the first
electron is called the first ionization energy.
Ionization Energy (cont.)
• Removing the second electron requires
more energy, and is called the second
ionization energy.
• Each successive ionization requires more
energy, but it is not a steady increase.
Ionization Energy (cont.)
• The ionization at which the large increase
in energy occurs is related to the number of
valence electrons.
• First ionization energy increases from left to
right across a period.
• First ionization energy decreases down a
group because atomic size increases and
less energy is required to remove an electron
farther from the nucleus.
Ionization Energy (cont.)
Ionization Energy (cont.)
Ionization Energy (cont.)
Ionization Energy (cont.)
Metallic Character
• Chemical properties associated with
elements classified as metals
• Examples??
• As you move across the periodic table
from left to right, metallic character
decreases
• As you move down a group, the metallic
character increases
Metallic Character
Nonmetallic Character
• Chemical properties
associated with chemicals
classified as nonmetals
• Examples???
• As you move across the
periodic table from left to
right, non metallic
character increases
• As you move down a
group, non metallic
character decreases
Electronegativity
• The measure of an
atom’s ability to attract
electrons
• As you move left to right
on the periodic table,
electronegativity
increases
• As you move down a
group decreases due to
the longer distance
between the outer
electrons and the nucleus
Ionization Energy (cont.)
• The electronegativity of an element
indicates its relative ability to attract
electrons in a chemical bond.
• Electronegativity decreases down a group
and increases left to right across a period.
Section 6.3 Assessment
The lowest ionization energy is the ____.
A. first
B. second
C. third
D
C
A
0%
B
D. fourth
A. A
B. B
C. C
0%
0%
0%
D. D
Section 6.3 Assessment
The ionic radius of a negative ion
becomes larger when:
A. moving up a group
B. moving right to left across period
D
A
0%
C
D. the ion loses electrons
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. moving down a group
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Concepts in Motion
Section 6.1 Development of the
Modern Periodic Table
Key Concepts
• The elements were first organized by increasing
atomic mass, which led to inconsistencies. Later,
they were organized by increasing atomic number.
• The periodic law states that when the elements are
arranged by increasing atomic number, there is a
periodic repetition of their chemical and physical
properties.
• The periodic table organizes the elements into periods
(rows) and groups (columns); elements with similar
properties are in the same group.
Section 6.1 Development of the
Modern Periodic Table (contd.)
Key Concepts
• Elements are classified as either metals, nonmetals,
or metalloids.
Section 6.2 Classification of
the Elements
Key Concepts
• The periodic table has four blocks (s, p, d, f).
• Elements within a group have similar chemical
properties.
• The group number for elements in groups 1 and 2
equals the element’s number of valence electrons.
• The energy level of an atom’s valence electrons equals
its period number.
Section 6.3 Periodic Trends
Key Concepts
• Atomic and ionic radii decrease from left to right
across a period, and increase as you move down a
group.
• Ionization energies generally increase from left to right
across a period, and decrease as you move down a
group.
• The octet rule states that atoms gain, lose, or share
electrons to acquire a full set of eight valence electrons.
• Electronegativity generally increases from left to right
across a period, and decreases as you move down a
group.
The actinide series is part of the
A. s-block elements.
B. inner transition metals.
C. non-metals.
D
C
A
0%
B
D. alkali metals.
A. A
B. B
C. C
0%
0%
0%
D. D
In their elemental state, which group has a
complete octet of valence electrons?
A. alkali metals
B. alkaline earth metals
D
A
0%
C
D. noble gases
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. halogens
Which block contains the transition
metals?
A. s-block
B. p-block
D
A
0%
C
D. f-block
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. d-block
An element with a full octet has how many
valence electrons?
A. two
B. six
D
A
0%
C
D. ten
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. eight
How many groups of elements are there?
A. 8
B. 16
C. 18
D
C
A
0%
B
D. 4
A. A
B. B
C. C
0%
0%
0%
D. D
Which group of elements are the least
reactive?
A. alkali metals
B. inner transition metals
D
A
0%
C
D. noble gases
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. halogens
On the modern periodic table, alkaline
earth metals are found only in ____.
A. group 1
B. s-block
D
A
0%
C
D. groups 13–18
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. p-block
Unreactive gases are mostly found where
on the periodic table?
A. halogens
B. group 1 and 2
D
A
0%
C
D. f-block
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. group 18
Bromine is a member of the
A. noble gases.
B. inner transition metals.
C. earth metals.
D
C
A
0%
B
D. halogens.
A. A
B. B
C. C
0%
0%
0%
D. D
How many groups does the d-block span?
A. two
B. six
C. ten
D
C
A
0%
B
D. fourteen
A. A
B. B
C. C
0%
0%
0%
D. D
Click on an image to enlarge.
Table 6.4
Noble Gas Electron
Configuration
Figure 6.5
The Periodic Table
Figure 6.11 Trends in Atomic Radii
Figure 6.18 Trends in Electronegativity
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