Download Periodic Trends Notes 14-15

Ch. 14: Chemical Periodicity
Standard: Matter consists of atoms that have internal structures that
dictate their chemical and physical behavior.
Targets:
• Describe the arrangement of elements in the periodic table in order of
increasing atomic number.
• Distinguish between the terms group and period.
• Apply the relationship between the electron arrangement of elements
and their position in the periodic table.
• Apply the relationship between the number of electrons in the highest
occupied energy level for an element and its position in the periodic
table.
• Discuss the similarities and differences in the chemical properties of
elements in the same group.
• Describe and explain the group and periodic trends in atomic radii, first
ionization energies and electronegativities.
Describe the arrangement of elements in the periodic
table in order of increasing atomic number.
Development of the Periodic Table
• Johan Dobereiner
Grouped similar elements into groups of 3
(triads) such as chlorine, bromine, and
iodine. (1817-1829).
• John Newlands
Found every eighth element (arranged by
atomic weight) showed similar properties.
Law of Octaves (1863).
• Dmitri Mendeleev
Arranged elements by similar properties
but left blanks for undiscovered elements
(1869).
Describe the arrangement of elements in the periodic
table in order of increasing atomic number.
Distinguish between the terms group and period.
Development of the Periodic Table
• Henry Mosley
Arranged the elements by increasing
atomic number instead of mass (1913)
• Glen Seaborg
Discovered the transuranium elements
(93-102) and added the actinide and
lanthanide series (1945)
Elements are arranged by increasing
atomic number into periods (rows)
and groups or families (columns)
Describe the arrangement of elements in the periodic
table in order of increasing atomic number.
Arrangement of the Periodic Table
• Metals
– Left side of the periodic table
(except hydrogen).
– High electrical conductivity,
high luster, ductile,
malleable
– Alkali metals: Group 1
– Alkaline earth metals: Group 2
– Transition metals: Group 3-12,
lanthanide & actinide series
Describe the arrangement of elements in the periodic
table in order of increasing atomic number.
Arrangement of the Periodic Table
• Nonmetals
– Right side of the periodic
table
– Poor conductors,
nonlustrous,
nonmalleable, dull
– Halogens: Group 17
– Noble gases: Group 18
•
Describe the arrangement of elements in the periodic
table in order of increasing atomic number.
Arrangement of the Periodic Table
• Metalloids
– Between metals and
nonmetals
– Properties
intermediate
between metals and
nonmetals
Apply the relationship between the electron arrangement
of elements and their position in the periodic table.
Arrangement of the Periodic Table: pg 392-393
• Noble Gases: Outermost s and p
sublevels are filled. Group 18
– Ending configuration is s2p6
(except He)
– Eight valence electrons
(except He)
– Row number equals highest
energy level
Apply the relationship between the electron arrangement
of elements and their position in the periodic table.
Arrangement of the Periodic Table: pg 392-393
• Representative Elements:
Outermost s and p sublevels are
partially filled.
– Group 1-2 and 13-18
– 1 (s1); 2 (s2); 13 (s2p1); 14
(s2p2)…
– Group number equals valence
electrons
– Row number equals highest
energy level
• Transition Metals
– Filling the d & f sublevels
– Go to Chemistry: periodicty Trends
Practice 1
Apply the relationship between the electron arrangement
of elements and their position in the periodic table.
Shortcut Electron Configuration
Based on the electron configuration of the noble gases.
He ends in 1s2; Ne ends in 2p6; Ar ends in 3p6; Kr ends in 4p6; etc.
•
Write the electron configuration and orbital filling diagram for Se
– Se has 34 electrons
– Go back to the previous noble gas: Ar (18 electrons). Begin the
configuration with [Ar] which accounts for 18 electrons and then
begin with 4s2. Continue until you reach 34 electrons
– [Ar]4s23d104p4
– [Ar] __ __ __ __ __ __ __ __ __
4s
3d
4p
Apply the relationship between the electron arrangement
of elements and their position in the periodic table.
Shortcut Electron Configuration
•
Write the electron configuration and orbital filling diagram for Au
– Au has 79 electrons
– Go back to the previous noble gas: Xe (54 electrons). Begin the
configuration with [Xe] which accounts for 54 electrons and
then begin with 6s2. Continue until you reach 79 electrons
– [Xe]6s24f145d9
– [Xe] __ __ __ __ __ __ __ __ __ __ __ __ __
6s
4f
5d
GO TO NOBLE GAS CONFIGURATION PRACTICE Handout
Apply the relationship between the number of electrons
in the highest occupied energy level for an element and
its position in the periodic table.
Shortcut Electron Configuration
Electron dot diagrams
Group 1: 1 dot
X
Group 15: 5 dots
X
Group 2: 2 dots
X
Group 16: 6 dots
X
Group 13: 3 dots
X
Group 17: 7 dots
X
Group 14: 4 dots
X
Group 18: 8 dots (except He)
X
GO TO WHITE BOARD LEWIS DOT PRACTICE and Noble gas
configuration practice.
Trends in the Periodic Table
• GO TO PERIOD TRENDS GRAPH
ACTIVITY
• GO TO REACTIVITY LAB
Periodic Trend Definitions
• Atomic Radius: half the internuclear distance
between two atoms of the same element (pm)
13
Periodic Trend Definitions
• Electronegativity: a measure of the tendency
of an atom in a molecule to attract a pair of
shared electrons towards itself
14
• trends are easier to understand if
you comprehend the following
• the ability of an atom to “hang on to”
or attract its valence electrons is the
result of two opposing forces
– the attraction between the electron and
the nucleus
– the repulsions between the electron in
question and all the other electrons in
the atom (often referred to the shielding
effect)
– the net resulting force of these two is
referred to effective nuclear charge
This is a simple, yet very good picture. Do you understand it?
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Atomic Radii
• The radius of an atom, measured in pm (picometers)
• Periodic trend (Period 3 Trend)
– Atomic size decreases as you move across a period.
– The increase in nuclear charge increases the attraction to the
outer shell so the outer energy level progressively becomes
closer to the nucleus
• Group trend for Alkali metals & Halogens
– Atomic size increases as you move down a group of the
periodic table.
– Adding higher energy levels
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Atomic Radii
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3.2.2 Describe and explain the trends in atomic
radii, electronegativity's and reactivity.
Electronegativity
Tendency for the atoms of the element to attract electrons
when they
are chemically combined with atoms of another element.
Helps predict the type of bonding (ionic/covalent).
• Periodic Trend (Period 3 Trend)
– Increases as you move from left to right across a period.
– Nonmetals have a greater attraction for electrons than metals & there is
a greater nuclear charge that can attract electrons
•
Group trend for Alkali metals & Halogens
– Generally decreases as you move down a group in the periodic table.
– For metals, the lower the inner nuclear force the more reactive.
– For nonmetals, the higher the inner nuclear force the more reactive.
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Electronegativity
22
3.2.2 Describe and explain the trends in atomic
radii, ionic radii, first ionization energies,
electronegativities and melting points
Reactivity
The relative capacity of an atom, molecule or radical to
undergo a chemical reaction with another atom, molecule or
radical.
• Don’t worry about the periodic trend!!!
•
Group trend for Alkali metals
– Increases as you move down group 1 in the periodic table
– Since alkali metals are more likely to lose an electron, the ones with the
lowest inner nuclear force are the most reactive since they require the
least amount of energy to lose a valence electron.
•
Group trend for Halogens
– Decreases as you move down group 17 in the periodic table
– Since halogens are more likely to gain an electron, the ones with the
greatest electronegativity are the most reactive since they are most
effective at gaining a valence electron.
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Discuss the similarities and differences in the
chemical properties of elements in the same
group.
Group 1: Alkali Metals
•
•
•
•
•
Have 1 valence electron
Shiny, silvery, soft metals
React with water & halogens
Oxidize easily (lose electrons)
Reactivity increases down the
group
Group 17: Halogens
•
•
•
•
Have 7 valence electrons
Colored gas (F2, Cl2); liquid (Br2);
Solid (I2)
Oxidizer (gain electrons)
Reactivity decreases down the
group