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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
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.
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Section 6-1
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
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Section 6-1
Development of the Periodic Table

• In the 1700s, Lavoisier compiled a list of all
  the known elements of the time.

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Section 6-1
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.

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Section 6-1
Development of the Periodic Table (cont.)

• Newlands noticed when the elements were arranged
  by increasing atomic mass, their properties
  repeated every eighth element.

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Section 6-1
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.

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Section 6-1
Development of the Periodic Table (cont.)
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Section 6-1
The Modern Periodic Table

• The modern periodic table contains boxes which
  contain the element's name, symbol, atomic
  number, and atomic mass.

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Section 6-1
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.

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Section 6-1
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.

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Section 6-1
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.

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Section 6-1
The Modern Periodic Table (cont.)

• Non-metals are elements that are generally gases
  or brittle, dull-looking solids, and poor
  conductors of heat and electricity.

• Group 17 is composed of highly reactive elements
  called halogens.
• Group 18 gases are extremely unreactive and
  commonly called noble gases.

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Section 6-1
The Modern Periodic Table (cont.)

• Metalloids have physical and chemical properties
  of both metals and non-metals, such as silicon
  and germanium.

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Section 6-1
The Modern Periodic Table (cont.)
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Section 6-1
Section 6.1 Assessment
What is a row of elements on the periodic table
called? A. octave B. period C. group
D. transition

1. A
2. B
3. C
4. D

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Section 6-1
Section 6.1 Assessment
What is silicon an example of? A. metal
B. non-metal C. inner transition metal
D. metalloid

1. A
2. B
3. C
4. D

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End of Section 6-1
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Section 6-2
Section 6.2 Classification of the Elements

• Explain why elements in the same group have
  similar properties.

valence electron electron in an atom's outermost
orbitals determines the chemical properties of
an atom

• Identify the four blocks of the periodic table
  based on their electron configuration.

Elements are organized into different blocks in
the periodic table according to their electron
configurations.
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Section 6-2
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.

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Section 6-2
Organizing the Elements by Electron Configuration
(cont.)

• The energy level of an elements 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.

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Section 6-2
Organizing the Elements by Electron Configuration
(cont.)
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Section 6-2
The s-, p-, d-, and f-Block Elements

• The shape of the periodic table becomes clear if
  it is divided into blocks representing the atoms
  energy sublevel being filled with valence
  electrons.

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Section 6-2
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.

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Section 6-2
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.

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Section 6-2
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.

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Section 6-2
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.

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Section 6-2
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 C. g-block D. f-block

1. A
2. B
3. C
4. D

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Section 6-2
Section 6.2 Assessment
Which block spans 14 elemental groups?
A. s-block B. p-block C. f-block D. g-block

1. A
2. B
3. C
4. D

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End of Section 6-2
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Section 6-3
Section 6.3 Periodic Trends

• Compare period and group trends of several
  properties.

principal energy level the major energy level of
an atom

• Relate period and group trends in atomic radii to
  electron configuration.

ion ionization energy octet rule electronegativity
Trends among elements in the periodic table
include their size and their ability to lose or
attract electrons
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Section 6-3
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.

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Section 6-3
Atomic Radius (cont.)

• For elements that occur as molecules, the atomic
  radius is half the distance between nuclei of
  identical atoms.

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Section 6-3
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.

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Section 6-3
Atomic Radius (cont.)
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Section 6-3
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.

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Section 6-3
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.

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Section 6-3
Ionic Radius (cont.)

• When atoms gain electrons, they can become
  larger, because the addition of an electron
  increases electrostatic repulsion.

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Section 6-3
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.

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Section 6-3
Ionic Radius (cont.)

• Both positive and negative ions increase in size
  moving down a group.

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Section 6-3
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.

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Section 6-3
Ionization Energy (cont.)
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Section 6-3
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.

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Section 6-3
Ionization Energy (cont.)
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Section 6-3
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.

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Section 6-3
Ionization Energy (cont.)
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Section 6-3
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.

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Section 6-3
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.

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Section 6-3
Ionization Energy (cont.)
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Section 6-3
Section 6.3 Assessment
The lowest ionization energy is the ____.
A. first B. second C. third D. fourth

1. A
2. B
3. C
4. D

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Section 6-3
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 C. moving down a group
D. the ion loses electrons

1. A
2. B
3. C
4. D

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End of Section 6-3
54
Resources Menu
Chemistry Online Study Guide Chapter
Assessment Standardized Test Practice Image
Bank Concepts in Motion
55
Study Guide 1
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.

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Study Guide 1
Section 6.1 Development of the Modern Periodic
Table (contd.)
Key Concepts

• Elements are classified as either metals,
  nonmetals, or metalloids.

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Study Guide 2
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 elements number of valence electrons.
• The energy level of an atoms valence electrons
  equals its period number.

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Study Guide 3
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.

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Chapter Assessment 1
The actinide series is part of the A. s-block
elements. B. inner transition metals. C. non-meta
ls. D. alkali metals.

1. A
2. B
3. C
4. D

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Chapter Assessment 2
In their elemental state, which group has a
complete octet of valence electrons? A. alkali
metals B. alkaline earth metals C. halogens
D. noble gases

1. A
2. B
3. C
4. D

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Chapter Assessment 3
Which block contains the transition metals?
A. s-block B. p-block C. d-block D. f-block

1. A
2. B
3. C
4. D

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Chapter Assessment 4
An element with a full octet has how many valence
electrons? A. two B. six C. eight D. ten

1. A
2. B
3. C
4. D

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Chapter Assessment 5
How many groups of elements are there? A. 8
B. 16 C. 18 D. 4

1. A
2. B
3. C
4. D

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STP 1
Which group of elements are the least reactive?
A. alkali metals B. inner transition metals
C. halogens D. noble gases

1. A
2. B
3. C
4. D

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STP 2
On the modern periodic table, alkaline earth
metals are found only in ____. A. group 1
B. s-block C. p-block D. groups 1318

1. A
2. B
3. C
4. D

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STP 3
Unreactive gases are mostly found where on the
periodic table? A. halogens B. group 1 and 2
C. group 18 D. f-block

1. A
2. B
3. C
4. D

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STP 4
Bromine is a member of the A. noble gases.
B. inner transition metals. C. earth metals.
D. halogens.

1. A
2. B
3. C
4. D

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STP 5
How many groups does the d-block span? A. two
B. six C. ten D. fourteen

1. A
2. B
3. C
4. D

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CIM
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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