Download Ch. 6 - The Periodic Table

Chapter 6
6.1

In a self-service store, the
products are grouped
according to similar
characteristics. With a logical
classification system, finding
and comparing products is
easy. You will learn how
elements are arranged in the
periodic table and what that
arrangement reveals about
the elements.
6.1

Searching For an Organizing Principle
◦ How did chemists begin to organize the known
elements?
 Chemists used the properties of elements to sort them
into groups (vertical columns).
6.1
 Chlorine, bromine, and iodine have very similar chemical
properties.
6.1
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Mendeleev’s Periodic Table
◦ How did Mendeleev organize his periodic table?
 Mendeleev organized elements into groups based on a
set of repeating properties and according to increasing
atomic mass.
 He used the periodic table to predict the properties of
undiscovered elements.
6.1
 An Early Version of Mendeleev’s Periodic Table
6.1
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The Periodic Law
◦ How is the modern periodic table organized?
 In the modern periodic table, elements are arranged in
order of increasing atomic number.
6.1
 The periodic law: When elements are arranged in order
of increasing atomic number, there is a periodic
repetition of their physical and chemical properties.
 The properties of the elements within a period (row)
change as you move across a period from left to right.
 This same pattern of properties then repeats across the
next period.
6.1
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Metals, Nonmetals, and Metalloids
◦ What are three broad classes of elements?
 Three classes of elements are metals, nonmetals, and
metalloids.
6.1
 Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
 Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
◦ Metals
 Metals are good conductors of heat and electric
current.
 80% of elements are metals.
 All metals are solids at room temperature except mercury,
which is a liquid.
 Metals have a high luster, are ductile, and are malleable.
6.1
 Uses of Iron, Copper, and Aluminum
6.1
 Uses of Iron, Copper, and Aluminum
6.1
 Uses of Iron, Copper, and Aluminum
6.1
 Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
◦ Nonmetals
 In general, nonmetals are poor conductors of heat and
electric current.
 Most nonmetals are gases at room temperature.
 A few nonmetals are solids, such as sulfur and
phosphorus.
 One nonmetal, bromine, is a dark-red liquid.
6.1
 Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
◦ Metalloids
 A metalloid generally has properties that are similar to
those of metals and nonmetals.
 The behavior of a metalloid can be controlled by changing
conditions.
 Metalloids are also known as semi-metals.
6.1
 If a small amount of boron is mixed with silicon, the
mixture is a good conductor of electric current. Silicon
can be cut into wafers, and used to make computer
chips.
◦ Across a period, the properties of elements become
less metallic and more nonmetallic.
◦ Down a group (column), the properties of elements
become more metallic and less nonmetallic.
Where is the most metallic element?
6.2

A coin may contain much
information in a small
space—its value, the year
it was minted, and its
country of origin. Each
square in a periodic table
also contains information.
You will learn what types
of information are usually
listed in a periodic table.
6.2

Squares in the Periodic Table
◦ What type of information can be displayed in a
periodic table?
 The periodic table displays the symbols and names of
the elements, along with information about the
structure of their atoms.
http://www.privatehand.com/flash/elements.html
6.2

Some element families have names:
◦ The Group 1 elements are called alkali metals.
◦ The Group 2 elements are called alkaline earth
metals.
◦ The nonmetals of Group 17 are called halogens.
◦ The nonmetal gases of group 18 are called noble
gases.
6.2
6.2

Electron Configurations in Groups
◦ How can elements be classified based on their
electron configurations?
 Elements can be sorted into groups based on their
electron configurations.
 Elements in the same family have the same outer electron
configuration = valence electrons.
6.2
◦ The Noble Gases
 The noble gases are the elements in Group 18 of the
periodic table; all noble gases have a full outer
electron energy level.
Helium (He)
Neon (Ne)
Argon (Ar)
Krypton (Kr)
2
2-8
2-8-8
2-8-18-8
6.2
◦ The alkali metals.
 In atoms of the Group 1 elements below, there is only
one electron in the highest occupied energy level; one
valence electron.
Lithium (Li)
Sodium (Na)
Potassium (K)
2-1
2-8-1
2-8-8-1
6.2
◦ The carbon family
 In atoms of the Group 14 elements below, there are
four valence electrons.
Carbon (C)
Silicon (Si)
Germanium (Ge)
2-4
2-8-4
2-8-18-4
6.2
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Transition Elements
 There are two types of transition elements—transition
metals and inner transition metals. They are classified
based on their electron configurations.
6.2
 In atoms of a transition metal, the d sublevel is filling
with electrons.
 In atoms of an inner transition metal, the f sublevel is
filling with electrons.
6.2
◦ Blocks of Elements
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The Elements: Forged in Stars - YouTube
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In the Earth’s Crust:
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Oxygen
Silicon
Aluminum
Iron
Calcium
Sodium
Potassium
Magnesium
Titanium
Hydrogen
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Dissolved in the Oceans:
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Chlorine
Sodium
Magnesium
Sulfur
Calcium
Potassium
Bromine
Carbon
Strontium
Boron
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In the Atmosphere:
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Nitrogen
Oxygen
Argon
Neon
Helium
Krypton
Hydrogen
Xenon
Radon
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In the Sun:
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Hydrogen
Helium
Oxygen
Carbon
Nitrogen
Silicon
Magnesium
Neon
Iron
Sulfur
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In your body:
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Oxygen
Carbon
Hydrogen
Nitrogen
Calcium
Phosphorus
Sulfur
Potassium
Sodium
Chlorine
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Soft, silver-grey metals.
Low melting and boiling points.
One valence electron.
Most reactive: not found uncombined in
nature.
◦ Obtained in the pure form by electrolysis of their
fused salts.

Potassium Video - The Periodic Table of
Videos - University of Nottingham
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Relatively soft, but harder than alkali metals.
Two valence electrons.
Although not as reactive as alkali metals, still
very reactive and not found in nature in the
elemental state.
◦ Obtained in the pure form through electrolysis of
their fused salts.
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Densities, melting and boiling points are
higher than respective alkali metals.
Radium Video - The Periodic Table of Videos
- University of Nottingham
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Most are ductile, malleable and good
conductors of heat and electricity.
Compounds of transition metals tend to have
color.
One or two valence electrons.
Obtained from mineral deposits (ores) in the
earth’s crust (smelting).
Precious metals are used for currency among
other things.
Darmstadtium Video - The Periodic Table of
Videos - University of Nottingham
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Nonmetals.
Very reactive; not found in nature
uncombined.
◦ Obtained from the electrolysis of their fused salts.
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Seven valence electrons.
Commercial applications include antibacterial
properties.
Chlorine Video - The Periodic Table of Videos
- University of Nottingham
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Full outer electron level.
Non-reactive (inert) gases.
◦ Can be forced to combine with fluorine.
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Commercial applications include colored
signs lit up as discharge tubes.
Incandescent light bulbs are filled with argon.
Radon Video - The Periodic Table of Videos University of Nottingham
6.3

Sodium chloride (table
salt) produced the
geometric pattern in the
photograph. Such a
pattern can be used to
calculate the position of
nuclei in a solid. You will
learn how properties
such as atomic size are
related to the location of
elements in the periodic
table.
6.3
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Trends in Atomic Size
◦ What are the trends among the elements for atomic
size?
 The atomic radius is one half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
6.3
◦ Group and Periodic Trends in Atomic Size
 In general, atomic size increases from top to bottom
within a group and decreases from left to right across a
period.
 Down a group, atomic size increases due to additional
energy levels.
 Across a period atomic size decreases due to increasing
nuclear charge.
6.3
6.3
6.3
 Some compounds are composed of particles called
ions.
 An ion is an atom or group of atoms that has a positive or
negative charge.
 A cation is an ion with a positive charge.
 An anion is an ion with a negative charge.
6.3
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Ions
◦ How do ions form?
6.3
Positive ions form when an atom loses electron(s).
6.3
Negative ions form when an atom gains electron(s).
6.3
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Trends in Ionization Energy
◦ What are the trends among the elements for first
ionization energy, ionic size, and electronegativity?
 The energy required to remove an electron from an
atom is called ionization energy.
 The energy required to remove the first electron from an
atom is called the first ionization energy.
 The energy required to remove a second electron is called
the second ionization energy.
6.3
◦ Group and Periodic Trends in Ionization Energy
 First ionization energy tends to decrease from top to
bottom within a group and increase from left to right
across a period.
 Down a group increasing levels of electrons shield the
effect of the nucleus therefore reducing energy needed to
remove an outer electron.
 Across a period there in no increase in energy levels, and
increasing nuclear charge makes it more difficult to
remove an outer electron.
6.3
6.3
6.3
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Trends in Ionic Size
◦ During reactions between metals and nonmetals,
metal atoms tend to lose electrons, and nonmetal
atoms tend to gain electrons. The transfer has a
predictable effect on the size of the ions that form.
 Cations are always smaller than the atoms from which
they form.
 Anions are always larger than the atoms from which
they form.
6.3
 Relative Sizes of Some Atoms and Ions
6.3
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Trends in Electronegativity
◦ Electronegativity is the ability of an atom to attract
electrons to itself when it is in involved in a bond.
 In general, electronegativity values decrease from top
to bottom within a group, and increase from left to
right across a period.
 Electronegativity decreases down a group because of
increasing atomic size and the shielding effect of inner
level electrons.
 Electronegativity increases across a period because of
decreasing atomic size and increasing nuclear charge.
6.3
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Summary of Trends
◦ What is the underlying cause of periodic trends?
 Periodic trends can be explained by variations in
atomic structure, nuclear charge, and shielding
effect.