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Chapter Presentation
Bellringer
Transparencies
Standardized Test Prep
Visual Concepts
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Chapter 4
The Periodic Table
Table of Contents
Section 1 How Are Elements Organized?
Section 2 Tour of the Periodic Table
Section 3 Trends in the Periodic Table
Section 4 Where Did the Elements Come From?
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Chapter 4
Section 1 How Are Elements
Organized?
Bellringer
• Make a list of things in the classroom that you think
are made from single elements. Make sure you think
about things you cannot see such as air.
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Chapter 4
Section 1 How Are Elements
Organized?
Objectives
• Describe the historical development of the periodic
table.
• Describe the organization of modern periodic table
according to the periodic law.
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Chapter 4
Section 1 How Are Elements
Organized?
The Periodic Law
• Mendeleev’s principle of chemical periodicity is known
as the periodic law, which states that when the
elements are arranged according to their atomic
numbers, elements with similar properties appear at
regular intervals.
Organization of the Periodic Table
• Elements in each column of the periodic table have
the same number of electrons in their outer energy
level.
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Chapter 4
Section 1 How Are Elements
Organized?
The Periodic Law, continued
Organization of the Periodic Table, continued
• The electrons in the outer shell are called valence
electrons.
• Valence electrons are found in the outermost shell
of an atom and that determines the atom’s chemical
properties.
• Elements with the same number of valence electrons
tend to react in similar ways.
• Because s and p electrons fill sequentially, the number
of valence electrons in s- and p-block elements are
predictable.
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Chapter 4
Section 1 How Are Elements
Organized?
Blocks of the Periodic Table
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Chapter 4
Section 1 How Are Elements
Organized?
The Periodic Law, continued
Organization of the Periodic Table, continued
• A vertical column on the periodic table is called a
group. Elements in a group share chemical properties.
• A horizontal row on the periodic table is called a
period. Elements in the same period have the same
number of occupied energy levels.
• Example: all elements in Period 2 have atoms
whose electrons occupy two principal energy levels,
including the 2s and 2p orbitals.
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Chapter 4
Section 1 How Are Elements
Organized?
The Periodic Law, continued
Organization of the Periodic Table, continued
• The periodic table
provides information
about each element.
• atomic number
• symbol
• name
• average atomic mass
• electron configuration
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Chapter 4
Visual Concepts
Periodic Table Overview
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HOMEWORK
SECTION REVIEW
Pg. 122
Q (5-14)
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Section 4.1 Review, pg.122
5. State the periodic law.
When elements are arranged according to their atomic
numbers, elements with similar properties appear at
regular intervals.
6. What do elements in the same period have in
common?
They have the same number of occupied energy levels
7. What do elements in the same group have in
common?
The same number of valence electrons
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8. Why can Period 1 contain a maximum of two
elements?
The first energy level of an atom can contain only two
electrons.
9. In which period and group is the element whose
electron configuration is [Kr]5s1?
Period 5, group 1
10. Write the outer electron configuration for the Group
2 element in Period 6.
6S2
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11. What determines the number of elements
found in each period in the periodic table?
The number of s and p electrons in that level + previous
level d- electrons + next previous level of f- electrons
12. Are elements with similar chemical properties
more likely to be found in the same period or in the same
group? Explain your answer
Same group, because they have same number of
valence electrons. Valence electrons determine the
chemical properties of the elements.
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13. How many valence electrons does phosphorus
have?
5
14. What would you expect the electron configuration of
element 113 to be?
[Rn] 7s2 5f14 6d10 7p1
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Chapter 4
Section 2 Tour of the Periodic Table
Bellringer
• On a blank periodic table, label each group by the
electron configuration of the valence electrons,
assuming the configuration follows the pattern given
by the aufbau principle.
• This pattern applies to all the main-group elements,
but there are many exceptions in the transition
metals in the center of the table.
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Chapter 4
Section 2 Tour of the Periodic Table
Objectives
• Locate the different families of main-group elements
on the periodic table, describe their characteristic
properties, and relate their properties to their electron
configuration.
• Locate metals on periodic table, describe their
characteristic properties, and relate their properties
to their electron configuration.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements
• Elements in groups 1, 2, and 13–18 are known as the
main-group elements. Main-group elements are in
the s- and p-blocks of the periodic table.
• The electron configurations of the elements in each
main group are regular and consistent: the elements
in each group have the same number of valence
electrons.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
• The main-group elements are sometimes called the
representative elements because they have a wide
range of properties.
• At room temperature and atmospheric pressure,
many are solids, while others are liquids or gases.
• About half of the main-group elements are metals.
• Many are extremely reactive, while several are
nonreactive.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
• The main-group elements silicon and oxygen account
for four of every five atoms found on or near Earth’s
surface.
• Four groups within the main-group elements have
special names. These groups are:
• alkali metals (Group 1)
• alkaline-earth metals (Group 2)
• halogens (Group 17)
• noble gases (Group 18)
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
• Main-group are highlighted in the groups on the left
and right sides of the periodic table.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
The Alkali Metals Make Up Group 1
• Elements in Group 1 are called alkali metals.
• lithium, sodium, potassium, rubidium, cesium, and
francium
• Alkali metals are so named because they are metals
that react with water to make alkaline solutions.
• Because the alkali metals have a single valence
electron, they are very reactive.
• In losing its one valence electron, potassium achieves a
stable electron configuration.
• Alkali metals are never found in nature as pure
elements but are found as compounds.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
The Alkali Metals Make Up Group 1, continued
Physical Properties of Alkali Metals
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
The Alkaline-Earth Metals Make Up Group 2
• Group 2 elements are called alkaline-earth metals.
• The alkaline-earth metals are slightly less reactive
than the alkali metals.
• They are usually found as compounds.
• The alkaline-earth metals have two valence electrons
and must lose both their valence electrons to get to a
stable electron configuration.
• It takes more energy to lose two electrons than it takes to
lose just the one electron that the alkali metals must give
up to become stable.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
The Halogens, Group 17, Are Highly Reactive
• Elements in Group 17 of the periodic table are called
the halogens.
• The halogens are the most reactive group of nonmetal
elements.
• When halogens react, they often gain the one electron needed
to have eight valence electrons, a filled outer energy level.
• Because the alkali metals have one valence electron,
they are ideally suited to react with the halogens.
• The halogens react with most metals to produce
salts.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
The Noble Gases, Group 18, Are Unreactive
• Group 18 elements are called the noble gases.
• The noble gas atoms have a full set of electrons in their
outermost energy level.
• The low reactivity of noble gases leads to some special
uses.
• The noble gases were once called inert gases because
they were thought to be completely unreactive.
• In 1962, chemists were able to get xenon to react, making the
compound XePtF6.
• In 1979, chemists were able to form the first xenon-carbon
bonds.
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Chapter 4
Section 2 Tour of the Periodic Table
The Main-Group Elements, continued
Hydrogen Is in a Class by Itself
• Hydrogen is the most common element in the universe.
• It is estimated that about three out of every four
atoms in the universe are hydrogen.
• Because it consists of just one proton and one electron,
hydrogen behaves unlike any other element.
• Hydrogen is in a class by itself in the periodic table.
• With its one electron, hydrogen can react with many
other elements, including oxygen.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals
• The majority of elements, including many main-group
ones, are metals.
• Metals are recognized by its shiny appearance, but
some nonmetal elements, plastics, and minerals are
also shiny.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
The regions highlighted in blue indicate the elements that
are metals.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Metals Share Many Properties
• All metals are excellent conductors of electricity.
• Electrical conductivity is the one property that
distinguishes metals from the nonmetal elements.
• Some metals, such as manganese, are brittle.
• Other metals, such as gold and copper, are ductile
and malleable.
• Ductile means that the metal can be squeezed
out into a wire.
• Malleable means that the metal can be
hammered or rolled into sheets.
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Chapter 4
Visual Concepts
Properties of Metals: Malleability and Ductility
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Transition Metals Occupy the Center of the Periodic
Table
• The transition metals constitute Groups 3 through 12
and are sometimes called the d-block elements
because of their position in the periodic table.
• A transition metal is one of the metals that can use
the inner shell before using the outer shell to bond.
• A transition metal may lose one, two, or even three
valence electrons depending on the element with which
it reacts.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Transition Metals Occupy the Center of the Periodic
Table, continued
• Generally, the transition metals are less reactive than
the alkali metals and the alkaline-earth metals are.
• Some transition metals are so unreactive that they
seldom form compounds with other elements.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Lanthanides and Actinides Fill f-orbitals
• The elements in the first of these rows are called the
lanthanides because their atomic numbers follow the
element lanthanum.
• A lanthanide is a member of the rare-earth series of elements,
whose atomic numbers range from 58 (cerium) to 71
(lutetium).
• Elements in the row below the lanthanides are called
actinides because they follow actinium.
• An actinide is any of the elements of the actinide series, which
have atomic numbers from 89 (actinium, Ac) through 103
(lawrencium, Lr).
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Lanthanides and Actinides Fill f-orbitals, continued
• As one moves left to right along these rows, electrons
are added to the 4f orbitals in the lanthanides and to
the 5f orbitals in the actinides.
• The lanthanides and actinides are sometimes called
the f-block of the periodic table.
• The actinides are unique in that their nuclear structures
are more important than their electron configurations.
• Because the nuclei of actinides are unstable and
spontaneously break apart, all actinides are
radioactive.
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Chapter 4
Section 2 Tour of the Periodic Table
Most Elements Are Metals, continued
Other Properties of Metals
• An alloy is a solid or liquid mixture of two or more
metals.
• The properties of an alloy are different from the
properties of the individual elements.
• Often these properties eliminate some
disadvantages of the pure metal.
• A common alloy is brass, a mixture of copper and zinc.
• Brass is harder than copper and more resistant to
corrosion.
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Chapter 4
Visual Concepts
Comparing Metals, Metalloids, and Nonmetals
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HOMEWORK
SECTION REVIEW
Pg. 131
Q(1-9)
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1. Which group of elements is the most unreactive?Why?
The noble gases ( group 18)because they have a full
valence level of electrons
2. Why do groups among the main-group elements
display similar chemical behavior?
They have the same number of valence electrons
3. What properties do the halogens have in common?
They have 7 valence electrons, highly reactive, react
with many metals, particularly alkali metals to form
salts
4. Why is hydrogen set apart by itself?
It has one proton and one electron.
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5. How do the valence electron configurations of the
alkali metals compare with each other?
They have the same number of valence electrons in their
outer s orbital.
6. Why are the alkaline-earth metals less reactive than
the alkali metals?
Alkaline earth metals must lose 2 electrons instead of one
to become stable.
7. In which groups of the periodic table do the
transition metals belong?
Groups 3 through 12
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8. Why are the nuclear structures of the actinides
more important than the electron configurations of the
actinides?
Their nuclei are unstable and spontaneously break apart,
making them radioactive
9. What is an alloy?
A mixture of a metal and one or more other elements.
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Chapter 4
Section 3 Trends in the Periodic
Table
Bellringer
• Draw atomic models of lithium, magnesium, and
fluorine.
• From the models, predict whether the ions of these
elements will be larger or smaller than the atoms. Be
sure to justify your predictions.
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Chapter 4
Section 3 Trends in the Periodic
Table
Objectives
• Describe periodic trends in ionization energy, and
relate them to the atomic structures of the elements.
• Describe periodic trends in atomic radius, and relate
them to the atomic structures of the elements.
• Describe periodic trends in electronegativity, and
relate them to the atomic structures of the elements.
• Describe periodic trends in ionic size, electron
affinity, and melting and boiling points, relate them
to the atomic structures of the elements.
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Chapter 4
Section 3 Trends in the Periodic
Table
Periodic Trends
• The arrangement of the periodic table reveals trends in
the properties of the elements.
• A trend is a predictable change in a particular direction.
• Understanding a trend among the elements enables
you to make predictions about the chemical behavior of
the elements.
• These trends in properties of the elements in a group or
period can be explained in terms of electron
configurations.
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Chapter 4
Section 3 Trends in the Periodic
Table
Ionization Energy
• The ionization energy is the energy required to
remove an electron from an atom or ion.
A + ionization energy  A + + e 
neutral atom
ion electron
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Chapter 4
Section 3 Trends in the Periodic
Table
Ionization Energy, continued
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The elements of Group 1, the Alkali metals, are:
symbol electron configuration
lithium
Li
1s22s1.
sodium
Na
1s2 2s22p6 3s1
potassium K
1s2 2s2 2p6 3s2 3p6 4s1
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Symbol
of the electron configuration of the ion
ion
lithium ion Li +
1s22s1 1s2
Sodium
ion
Na +
1s2 2s22p6 3s1 1s2 2s22p6
potassiu
m ion
K+
1s2 2s2 2p6 3s2 3p6 4s11s2 2s2 2p6 3s2 3p6
Least Energy
most ionization energy
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less Energy
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Chapter 4
Section 3 Trends in the Periodic
Table
Ionization Energy, continued
Ionization Energy Decreases as You Move Down a
Group
• Each element has more occupied energy levels than
the one above it has.
• The outermost electrons are farthest from the nucleus in
elements near the bottom of a group.
• As you move down a group, each successive element
contains more electrons in the energy levels between
the nucleus and the outermost electrons.
• Electron shielding is the reduction of the attractive force
between a positively charged nucleus and its outermost
electrons due to the cancellation of some of the positive
charge by the negative charges of the inner electrons.
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Chapter 4
Section 3 Trends in the Periodic
Table
Ionization Energy, continued
Ionization Energy Decreases as You Move Down a
Group, continued
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Chapter 4
Section 3 Trends in the Periodic
Table
Ionization Energy, continued
Ionization Energy Increases as You Move Across a
Period
• Ionization energy tends to increase as you move from
left to right across a period.
• From one element to the next in a period, the number
of protons and the number of electrons increase by one
each.
• The additional proton increases the nuclear charge.
• A higher nuclear charge more strongly attracts the outer
electrons in the same energy level, but the electron-shielding
effect from inner-level electrons remains the same.
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Write the electron configuration for each of the following
elements.
Sodium Na
1s2 2s22p6 3s1
Least ionization Energy
Magnesium Mg
1s2 2s22p6 3s2
More ionization E
Aluminum Al
1s2 2s22p6 3s2 3p1
More E
Silicon Si
1s2 2s22p6 3s2 3p2
Most E
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Chapter 4
Visual Concepts
Ionization
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Chapter 4
Section 3 Trends in the Periodic
Table
Atomic Radius
• The exact size of an atom is hard to determine.
• The volume the electrons occupy is thought of as an
electron cloud, with no clear-cut edge.
• In addition, the physical and chemical state of an atom
can change the size of an electron cloud.
• One method for calculating the size of an atom involves
calculating the bond radius, which is half the distance
from center to center of two like atoms that are bonded
together.
• The bond radius can change slightly depending on
what atoms are involved.
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Chapter 4
Section 3 Trends in the Periodic
Table
Atomic Radius, continued
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Chapter 4
Visual Concepts
Bond Length
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Chapter 4
Section 3 Trends in the Periodic
Table
Atomic Radius, continued
Atomic Radius Increases as You Move Down a Group
• As you proceed from one element down to the next in a
group, another principal energy level is filled.
• The addition of another level of electrons increases the
size, or atomic radius, of an atom.
• Because of electron shielding, the effective nuclear
charge acting on the outer electrons is almost constant
as you move down a group, regardless of the energy
level in which the outer electrons are located.
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Chapter 4
Section 3 Trends in the Periodic
Table
Atomic Radius, continued
Atomic Radius Decreases as You Move Across a
Period
• As you move from left to right across a period, each
atom has one more proton and one more electron than
the atom before it has.
• All additional electrons go into the same principal
energy level—no electrons are being added to the
inner levels.
• Electron shielding does not play a role as you move across a
period.
• As the nuclear charge increases across a period, the
effective nuclear charge acting on the outer electrons
also increases.
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Chapter 4
Section 3 Trends in the Periodic
Table
Periodic Trends of Radii
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity
• Not all atoms in a compound share electrons equally.
• Knowing how strongly each atom attracts bonding
electrons can help explain the physical and chemical
properties of the compound.
• Linus Pauling, an American chemists, made a scale of
numerical values that reflect how much an atom in a
molecule attracts electrons, called electronegativity
values.
• Electronegativity is a measure of the ability of an
atom in a chemical compound to attract electrons.
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Electronegativity, continued
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity, continued
• The atom with the higher electronegativity will pull on
the electrons more strongly than the other atom will.
• Fluorine is the element whose atoms most strongly
attract shared electrons in a compound. Pauling
arbitrarily gave fluorine an electronegativity value of
4.0.
• Values for the other elements were calculated in
relation to this value.
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity, continued
Electronegativity Decreases as You Move Down a
Group
• Electronegativity values generally decrease as you
move down a group.
• The more protons an atom has, the more strongly it
should attract an electron.
• However, electron shielding plays a role again.
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity, continued
Electronegativity Increases as You Move Across a
Period
• Electronegativity usually increases as you move left to
right across a period.
• As you proceed across a period, each atom has one
more proton and one more electron—in the same
principal energy level—than the atom before it has.
• Electron shielding does not change as you move
across a period because no electrons are being
added to the inner levels.
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity, continued
Electronegativity Increases as You Move Across a
Period, continued
• The effective nuclear charge increases across a period.
• As this increases, electrons are attracted much
more strongly, resulting in an increase in
electronegativity.
• The increase in electronegativity across a period is
much more dramatic than the decrease in
electronegativity down a group.
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Chapter 4
Section 3 Trends in the Periodic
Table
Electronegativity, continued
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends
• The effective nuclear charge and electron shielding are
often used in explaining the reasons for periodic trends.
• Effective nuclear charge and electron shielding also
account for two other periodic trends–ionic size and
electron affinity.
• The trends in melting and boiling points are determined
by how electrons form pairs as d orbitals fill.
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Ionic Size and Electron Affinity
• Like atomic size, ionic size has periodic trends.
• As you proceed down a group, the outermost electrons
in ions are in higher energy levels.
• The ionic radius usually increases as you move
down a group.
• This trends hold for both positive and negative ions.
• Metals tend to lose one or more electrons and form a
positive ion.
• As you move across a period, the ionic radii of
metal cations tend to decrease because of the
increasing nuclear charge.
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Ionic Size and Electron Affinity,
continued
• The atoms of nonmetal elements in a period tend to
gain electrons and form negative ions.
• As you proceed through the anions on the right of a
period, ionic radii still tend to decrease because of the
anions’ increasing nuclear charge.
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Ionic Size and Electron Affinity
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Ionic Size and Electron Affinity,
continued
• The energy change that occurs when a neutral atom
gains an electron is called the atom’s electron affinity.
• This property of an atom is different from electronegativity.
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• The electron affinity tends to decrease as you move
down a group because of the increasing effect of
electron shielding.
• Electron affinity tends to increase as you move across a
period because of the increasing nuclear charge.
Ionization energies are always concerned with the
formation of positive ions. Electron affinities are the
negative ion equivalent.
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Ionic Size and Electron Affinity
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Chapter 4
Visual Concepts
Electron Affinity
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Chapter 4
Section 3 Trends in the Periodic
Table
Other Periodic Trends, continued
Periodic Trends in Melting and Boiling Points
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Visual Concepts
Melting Point
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Visual Concepts
Boiling Point
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Chapter 4
Section 3 Trends in the Periodic
Table
Additional Periodic Trends
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Chapter 4
Section 3 Trends in the Periodic
Table
Additional Periodic Trends, continued
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Chapter 4
Section 3 Trends in the Periodic
Table
Additional Periodic Trends, continued
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HOMEWORK
SECTION REVIEW
PG. 141
Q 111
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1. What is ionization energy?
Is the amount of energy needed to remove an
electron from an atom.
2. Why is measuring the size of an atom difficult?
An atom has an electron cloud that has no definite
boundary.
3. What can you tell about an atom that has high
electronegativity?
It will strongly attract other electrons in a compound.
4. How does electron shielding affect atomic size as
you move down a group?
Electron shielding contributes to the increase in atomic
size as you move down a group.
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5. What periodic trends exist for ionization energy?
Ionization energy decreases down a group and increases
from left to right across a period.
6. Describe one way in which atomic radius is defined.
Half the distance between the nuclei of two bonded
atoms.
7. Explain how the trends in melting and boiling points
differ from the other periodic trends.
They are not consistent. Across a period they may
increase, then decrease, then repeat the pattern.
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8. Why do both atomic size and ionic size increase as you
move down a group?
Another principal energy level is added as you move from
one element to the next, resulting in an increase in size.
9. How is electron affinity different from electronegativity?
Electron affinity is the attraction an atom has when it is not
bonded. Electronegativity is the attraction an atom has when
it is bonded to another atom.
10. What periodic trends exist for electronegativity?
Decreases down a group and increases from left to right
11. Why is electron shielding not a factor when you
examine a trend across a period?
Each atom across a period has the same number of energy
levels and therefore the same amount of shielding.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Bellringer
• Find technetium, promethium, and neptunium on a
blank periodic table.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Objectives
• Describe how the naturally occurring elements form.
• Explain how a transmutation changes one element
into another.
• Describe how particle accelerators are used to
create synthetic elements.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements
• Of all the elements listed in the periodic table, 93 are
found in nature.
• Three of these elements, technetium, Tc, promethium,
Pm, and neptunium, Np, are not found on Earth but
have been detected in the spectra of stars.
• Most of the atoms in living things come from just six
elements.
• carbon, hydrogen, oxygen, nitrogen, phosphorus,
and sulfur
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements, continued
Hydrogen and Helium Formed After the Big Bang
• Much of the evidence about the universe’s origin points
toward a single event: an explosion of unbelievable
violence, before which all matter in the universe could
fit on a pinhead. This event is known as the big bang.
• As the universe expanded, it cooled and some of the
energy was converted into matter in the form of
electrons, protons, and neutrons.
• As the universe continued to cool, these particles
started to join and formed hydrogen and helium
atoms.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements, continued
Hydrogen and Helium Formed After the Big Bang,
continued
• Gravity pulled these clouds of hydrogen closer and
closer.
• As the clouds grew more dense, pressures and
temperatures at the centers of the hydrogen clouds
increased, and stars were born.
• In the centers of stars, nuclear reactions took place.
• A nuclear reaction is a reaction that affects the
nucleus of an atom.
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Chapter 4
Visual Concepts
Nuclear Reaction
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Chapter 4
Visual Concepts
Nuclear Fusion
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements, continued
Other Elements Form by Nuclear Reactions in Stars
• Einstein’s equation E = mc2 describes the mass-energy
relationship quantitatively.
• Einstein’s equation shows that fusion reactions release very
large amounts of energy.
• The energy released by a fusion reaction is so great it
keeps the centers of the stars at very high
temperatures.
• During fusion in stars, two helium nuclei fuse to form a
beryllium nucleus, and gamma radiation is released.
• Repeated fusion reactions can form atoms as massive
as iron and nickel.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements, continued
Other Elements Form by Nuclear Reactions in Stars,
continued
• Very massive stars (stars whose masses are more than
100 times the mass of our sun) are the source of
heavier elements.
• When such a star has converted almost all of its
core hydrogen and helium into the heavier
elements up to iron, the star collapses and then
blows apart in an explosion called a supernova.
• All of the elements heavier than iron on the periodic
table are formed in this explosion.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Natural Elements, continued
Other Elements Form by Nuclear Reactions in Stars
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Chapter 4
Section 4 Where Did the Elements
Come From?
Nuclear Fusion: Stellar Formation of Carbon-12
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Chapter 4
Section 4 Where Did the Elements
Come From?
Transmutations
• In the Middle Ages, many early chemists tried to
change, or transmute, ordinary metals into gold.
• These early chemists did not realize that a
transmutation, whereby one element changes into
another, is a nuclear reaction. It changes the nucleus of
an atom and therefore cannot be achieved by ordinary
chemical means.
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Chapter 4
Visual Concepts
Transmutation
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Chapter 4
Section 4 Where Did the Elements
Come From?
Transmutations, continued
Transmutations Are a Type of Nuclear Reaction
• Because of the results of his experiments, Ernest
Rutherford believed that the nuclei in air had
disintegrated into the nuclei of hydrogen (protons) plus
the nuclei of some other atom.
• W. D. Harkins and P.M.S. Blackett studied this strange
phenomenon further.
• They concluded that the Y formed when an alpha
particle collided with a nitrogen atom in air to
produce an oxygen atom and a proton, and that a
transmutation had thereby occurred.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Synthetic Elements
• Chemists have synthesized, or created, more elements
than the 93 that occur naturally.
• These are synthetic elements.
• All of the transuranium elements, or those with more
than 92 protons in their nuclei, are synthetic
elements. To make them, one must use special
equipment, called particle accelerators.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Synthetic Elements, continued
The Cyclotron Accelerates Charged Particles
• Many of the first synthetic elements were made with the
help of a cyclotron, a particle accelerator, in which
charged particles are given one pulse of energy after
another, speeding them to very high energies.
• The particles then collide and fuse with atomic nuclei to
produce synthetic elements.
• There is a limit to the energies that can be reached with
a cyclotron and therefore a limit to the synthetic
elements that it can make.
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Chapter 4
Section 4 Where Did the Elements
Come From?
Synthetic Elements, continued
The Cyclotron Accelerates Charged Particles
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Chapter 4
Section 4 Where Did the Elements
Come From?
Synthetic Elements, continued
The Synchrotron Is Used to Create Superheavy
Elements
• Once the particles have been accelerated, they are
made to collide with one another to make superheavy
elements, which have atomic numbers greater than
106.
• Most superheavy elements exist for only a tiny fraction
of a second.
• Thirty seconds is a very long life span for a
superheavy element.
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Chapter 4
Standardized Test Preparation
Understanding Concepts
1. Which of the following elements is formed in stars?
A. curium
B. einsteinium
C. gold
D. mendelevium
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Chapter 4
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Understanding Concepts
1. Which of the following elements is formed in stars?
A. curium
B. einsteinium
C. gold
D. mendelevium
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Chapter 4
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Understanding Concepts
2. Why are the Group 17 elements, the halogens, the
most reactive of the nonmetal elements?
F. They have the largest atomic radii.
G. They have the highest ionization energies.
H. They are the farthest right on the periodic table.
I. They require only one electron to fill their outer
energy level.
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Chapter 4
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Understanding Concepts
2. Why are the Group 17 elements, the halogens, the
most reactive of the nonmetal elements?
F. They have the largest atomic radii.
G. They have the highest ionization energies.
H. They are the farthest right on the periodic table.
I. They require only one electron to fill their outer
energy level.
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Understanding Concepts
3. Which of the following is a property of noble gases as
a result of their stable electron configuration?
A. large atomic radii
B. high electron affinities
C. high ionization energies
D. a tendency to form both cations and anions
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3. Which of the following is a property of noble gases as
a result of their stable electron configuration?
A. large atomic radii
B. high electron affinities
C. high ionization energies
D. a tendency to form both cations and anions
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Understanding Concepts
4. Which of these is a transition element?
F. Ba
G. C
H. Fe
I. Xe
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Understanding Concepts
4. Which of these is a transition element?
F. Ba
G. C
H. Fe
I. Xe
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Understanding Concepts
5. How did the discovery of the elements that filled the
gaps in Mendeleev's periodic table increase
confidence in the periodic table?
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Understanding Concepts
5. How did the discovery of the elements that filled the
gaps in Mendeleev's periodic table increase
confidence in the periodic table?
Answer: The gaps were significant because they
predicted the properties of new elements that would
be discovered. Their discovery demonstrated that the
table was a useful tool for organizing information
about atoms.
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Understanding Concepts
6. Why is iodine placed after tellurium on the periodic
table if the atomic mass of tellurium is less than that
of iodine?
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Understanding Concepts
6. Why is iodine placed after tellurium on the periodic
table if the atomic mass of tellurium is less than that
of iodine?
Answer: Because the periodic table is based on atomic
number, not atomic mass. The atomic number of
iodine is one higher than the atomic number of
tellurium.
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Understanding Concepts
7. What is the outermost occupied energy level in atoms
of the elements in Period 4?
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Understanding Concepts
7. What is the outermost occupied energy level in atoms
of the elements in Period 4?
Answer: Level 4
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Chapter 4
Standardized Test Preparation
Reading Skills
Read the passage below. Then answer the questions.
The atomic number of beryllium is one less than
that of boron, which follows it on the periodic table.
Strontium, which is directly below beryllium in period 5
of the periodic table has 34 more protons and 34 more
electrons than beryllium. However, the properties of
beryllium resemble the much larger strontium more
than those of similar-sized boron.
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Chapter 4
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Reading Skills
8. The properties of beryllium are more similar to those
of strontium than those of boron because
A.
B.
C.
D.
strontium is larger than boron.
strontium and beryllium are both metals.
strontium has more electrons than boron.
strontium has the same number of valence
electrons as beryllium.
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Chapter 4
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Reading Skills
8. The properties of beryllium are more similar to those
of strontium than those of boron because
A.
B.
C.
D.
strontium is larger than boron.
strontium and beryllium are both metals.
strontium has more electrons than boron.
strontium has the same number of valence
electrons as beryllium.
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Reading Skills
9. Beryllium and strontium are both located in the
second column of the periodic table. To which of
these classifications do they belong?
F.
G.
H.
I.
alkali metals
alkaline earth metals
rare earth metals
transition metals
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9. Beryllium and strontium are both located in the
second column of the periodic table. To which of
these classifications do they belong?
F.
G.
H.
I.
alkali metals
alkaline earth metals
rare earth metals
transition metals
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10. Why is it easier to determine to which column of the
periodic table an element belongs than to determine
to which row it belongs, based on observations of its
properties?
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10. Why is it easier to determine to which column of the
periodic table an element belongs than to determine
to which row it belongs, based on observations of its
properties?
Answer: It is easier to determine the column because
all the elements in a column have the same outer
electron structure and, therefore, similar properties.
Properties of elements across a row of the table vary
widely.
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Interpreting Graphics
Use the diagram below to answer question 11.
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Chapter 4
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Interpreting Graphics
11. What process is represented by this illustration?
A. chemical reaction
B. ionization
C. nuclear fission
D. nuclear fusion
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Chapter 4
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Interpreting Graphics
11. What process is represented by this illustration?
A. chemical reaction
B. ionization
C. nuclear fission
D. nuclear fusion
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Chapter 4
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Interpreting Graphics
The graph below shows the ionization energies
(kilojoules per mole) of main-block elements. Use it to
answer questions 12 and 13.
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Interpreting Graphics
12. Which of these elements requires the most energy to
remove an electron?
F. argon
G. fluorine
H. nitrogen
I. oxygen
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Interpreting Graphics
12. Which of these elements requires the most energy to
remove an electron?
F. argon
G. fluorine
H. nitrogen
I. oxygen
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Reading Skills
13. Explain the trend in ionization energy within a group
on the periodic table.
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Chapter 4
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Reading Skills
13. Explain the trend in ionization energy within a group
on the periodic table.
Answer: Ionization energy tends to increase from left to
right across the table because elements have
increasingly more protons so the attraction on the
outer electrons is stronger.
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