Transcript Families of elements

Chapter 6 The Periodic Table
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Development/History of the
Modern Periodic Table
Using the Periodic Table
An Introduction to the Elements
Periodic Trends
Periodic Table
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Why Periodic????
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The properties of the elements repeat in in a
periodic way.
Invaluable tool for chemistry
Used for organization
History of the Periodic Table
 Timeline
 Trace
the development of the Periodic
Table by making a timeline
Aristotle
 Newlands
 Dobereiner
 Meyer
 Mendeleev
 Moseley
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The Basics
 Elements
are
arranged by
atomic number
 Typical box
contains:
 Name
of the
element
 Symbol
 Atomic number
 Atomic mass
Periods Horizontal Rows
Numbered 1-7
Groups: Vertical Columns
Numbered 1-18
Interactive Periodic Tables
 http://periodic.lanl.gov/index.shtml
 www.webelements.com
 www.chemicool.com
 http://education.jlab.org/itselemental/ele
016.html
The Families of Elements
http://www.privatehand.com/flash/elements.html
Classification of the Elements
Metals
Nonmetals
Metalloids
Metals
•Occupy
the left side
of the periodic table
•Have luster, shiny
•Solids at room temp
except Hg
•Ductile: ability to be
drawn into wires
•Malleable: ability to
be hammered into
sheets
•Excellent
conductors of heat
and electricity
•Tend to form
positive ions
NonMetals
 Occupy
the right side of the
Periodic Table
 Generally gases or brittle
solids
 Dull-looking
 Brittle
 Poor conductors of heat and
electricity
 Bromine is the only liquid
at room temp
 Tend to form negative ions
Metalloids
Characteristics of
metals and nonmetals
Classification of the Elements
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Families of elements share the same ending
electron configuration
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therefore they share similar chemical
characteristics
Valence Electrons: electrons in the highest
principal energy level
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Determine Chemical reactivity
Elements in a group share the same number of
valence electrons
The s, p, d and f blocks
Number of Valence Electrons
Elements on the right
• Nonmetals
• 4 or more valence electrons
• tend to gain electrons
• become negative ions
Elements on the left
• Metals
• 3 or less valence electrons
• tend to lose valence electrons
• form positive ions
Most Common Ions
Families of elements
 Elements
of the same family (group)
share structural and chemical
(behavioral) characteristics
 Alkali
Metals
 Alkaline Earth Metals
 Transition Elements
 Halogens
 Nobel Gases
Group 1: Alkali Metals
 Soft,
highly reactive
metals
 Usually stored under
oil or kerosene to
prevent their
interaction with air
and water
Properties of Alkali Metals
 React
vigorously with water
 Oxidize readily in air
 Good conductors of electricity
Alkali Metals
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Have one valence electron
 Will lose this electron very
easily
 when electron is lost the
metal gains a stable nonreactive noble gas
configuration
Comparison of the Reactivity of the Alkali Metals
http://www.youtube.com/watch?v=uixxJtJPVXk
Group 2: Alkaline Earth Metals
 Harder,
denser, stronger, and have
higher melting points than alkali metals
 All are reactive not as reactive as group 1
Alkaline Earth Metals
 Must
lose two electrons to gain a stable
configuration
Groups 3-12:Transition Metals
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Not as reactive as Groups 1 and 2
 Huge variety but all shiny
 Multi valent…form multiple ions
 d-block elements
 Also include: Inner Transition Elements (Rare
Earth Elements)
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Elements 58-71 Lanthanides
Elements 90-103 Actinides
Group 17: Halogens
 Most
reactive non-metals
 Combine easily with metals; especially
the alkali metals
Halogens
7
valence electrons, one short of a
stable octet.
 Will gain one electron to become stable
 -1 ions
Reaction of chlorine (a halogen) with sodium (an alkali
metal)
https://www.youtube.com/watch?v=1xT4OFS03jE
Element Dating
Hydrogen
 Most
common element in the universe
 Chemical family by itself because it
behaves so differently
 Reacts with most other elements
 Rarely found in a free state in nature
 1 valence electron
The Hindenberg
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Filled with H
 Very reactive with
oxygen gas
 He used in blimps
today
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much less reactive
than H
Group 18: Noble Gases
 Very
low reactivity
 Filled valence shells: s and p levels in
the highest principal energy levels are
full
 Very stable electron configuration
 Many uses: signs, weather balloons and
the airships (Blimps)
The Octet Rule
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Atoms tend to gain, lose or share electrons in
order to acquire a full set of eight valence
electrons.
Elements on the left (metals) tend to lose
valence electrons and form positive ions
 Elements on the right (nonmetals) tend to
gain electrons to become negative ions
Periodic Trends
 Properties
of Elements tend to occur in
a predictable way
 Known as a trend, as you move across
a period or down a group
 Knowing element trends allows us to
make predictions about an element’s
behavior
Periodic Properties
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Properties
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Atomic Radius
Ionic Radius
Electronegativity
Ionization Energy
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Questions we will
answer:
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Definition
How does the
property vary across
the table?
Why?
How does it vary
down a group?
Why?
Atomic Radius
• For elements that occur as molecules, the
atomic radius is half the distance between
nuclei of identical atoms.
Atomic Radius
 The
atomic radius is a measure of
the size of an atom.
 The larger the radius, the larger is
the atom.
Trends in Atomic Radius
 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.
Trends in Atomic Radius
 The
atomic radius decreases as
you move across a period
 Why?
 Increased nuclear charge pulls the
electrons in tighter
 Added electrons are in the same
principal energy levels
Group Trends in Atomic Radius
 Atomic
Radius increases as you move
down a group
 Why?
The increasing number of electrons are in
higher energy levels and instead of
pulling the electrons closer to the
nucleus we see the …
Atomic Radius
• Atomic radius generally increases as you
move down a group.
• The outermost orbital size increases down a
group, making the atom larger.
Shielding Effect
 More
inner electrons shield the
outer electron from the nucleus and
reduce their attraction to the
nucleus therefore the overall atomic
radius is larger
Ionic Radius
 Atoms
can gain or lose electrons to
form ions
 Ion:
an atom with a charge
 Recall
that atoms are neutral in charge,
 If an electron is lost, then the overall
charge is positive
 If an electron is gained the atom
becomes negative
Positive Ion (Cation) Formation
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When atoms lose electrons
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Radius always becomes smaller
Because…
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The loss of a valence electron can leave an
empty outer orbital resulting in a small radius.
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Electrostatic repulsion decreases allowing the
electrons to be pulled closer to the radius.
Negative Ion (Anion) Formation
 When
atoms gain electrons
 Radius always increases
Why?
 More
electrons mean more electrostatic
repulsion resulting in increased diameter.
Period Trend for Ionic Radius
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As you move left to
right across a period
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the ionic radius gets
smaller for the
positive ions
The ionic radius for
the negative ions
also decreases
Group Trend for Ionic Radius
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Both positive and
negative ions
increase in size
moving down a
group.
Ionic Radius
Ionization Energy
 the
amount of energy need to remove
an electron from a specific atom or ion
in its ground state in the gas phase
 High
Ionization Energy: atom is holding
onto electrons very strongly
 Low Ionization Energy: atom is holding
electrons less tightly
 For
any element (A) the process of
removing an electron can be
represented as follows:
 A + energy -----> A+ + e What is the periodic trend in
ionization energy? Why?
Trends for
Ionization Energy
 Generally
increases as you move
across a period
 because
increased nuclear charge causes
an increased hold on the electrons
 Ionization
Energy decreases as you
move down a group
 due
to increasing atomic size
Successive Ionization Energies
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There is an ionization energy for each
electron that is removed from an atom
 After the valence electrons are removed
Ionization Energies Jump Dramatically
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First Ionization Energy: removes 1 electron
Second Ionization Energy: removes a second
electron
Third Ionization Energy: removes a third electron
Comparing Successive Ionization
Energies
Trends in Ionization Energy
Electronegativity
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The ability of an an atom to attract electrons
to itself when it is combined with another
atom
 Expressed in terms of a relative scale:
fluorine is assigned a value of 4 and all other
elements are calculated relative to this.
 The units of electronegativity are arbitrary
units called Paulings.
 Noble gases have no values because of few
chemical compounds
Electronegativity
 Greater
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the electronegativity
the higher an atom’s ability to pull an
electron to itself when it is bonded to
another atom
 What
are the periodic trends in
electronegativity?
 Why?
Trends in Electronegativity
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Electronegativity Increases as you move
across a period
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Electronegativity decreases you move down
a group
Where are the elements with highest
electronegativity?
Where are the elements with lowest
electronegativity?
Electronegativity
Summary of Trends
Another Summary