Transcript The Modern Periodic Table Chapter 6

The Modern Periodic Table
Chapter 6
Objectives for Friday:
4. Define period, group, and family.
 5. Explain the different systems for
numbering the groups on the Periodic
table.
 6. Name and locate on the periodic table:
metals, nonmetals, metalloids, noble
gases, alkali metals, alkaline earth metals,
halogens, chalcogens, Lanthanides,
Actinides, Transition metals, and Inner
transition metals.
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Objectives:
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7. Apply the octet rule in predicting the stability of an
element.
8. Predict the electron configuration of the outer energy
level of any element from its position on the Periodic
table.
9. Recognize exceptions to the Aufbau order due to halffull or full d-sublevels.
10. Define atomic radius, ionic radius, ionization energy,
electron affinity, and electronegativity.
11. Predict trends in atomic radius, ionic radius,
oxidation number, ionization energy, electron affinity, and
electronegativity.
12. Explain trends in terms of nuclear charge and
shielding effects.
Arrangement and Nomenclature
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Rows are called periods
Columns are designated as groups
Each column in the main table and each row
at the bottom is also designated an individual
family
Groups 1A, 2A, and 3-8A are the main
groups, or representative elements
Groups 1B-8B are called the transition
elements
The Periodic Table With Atomic Symbols, Atomic
Numbers, and Partial Electron Configurations
Broad Periodic Table
Classifications
Representative Elements (main group):
filling s and p orbitals (Na, Al, Ne, O)
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Transition Elements: filling d orbitals
(Fe, Co, Ni)
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Lanthanide and Actinide Series (inner
transition elements): filling 4f and 5f
orbitals (Eu, Am, Es)
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Information Contained in the
Periodic Table
Each group member has the
same valence electron
configuration (these electrons
primarily determine an atom’s
chemistry).
2. The electron configuration of
any representative element.
1.
Information Contained in the
Periodic Table
Certain groups have special
names (alkali metals, alkaline
earth metals, chalcogens,
halogens, etc).
4. Metals and nonmetals are
characterized by their chemical
and physical properties.
3.
Special Names for
Groups in the
Periodic Table
Metals
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Metals makeup more than 75% of the
elements in the periodic table. Metals are
characterized by the following physical
properties:
– They have metallic shine or luster.
– They are usually solids at room temperature.
– They are malleable. Malleable means that metals
can be hammered, pounded, or pressed into
different shapes without breaking.
– They are ductile meaning that they can be drawn
into thin sheets or wires without breaking.
– They are good conductors of heat and electricity.
Metals (cont)
All B and most A elements are metals.
The B  At stairstep designates the border
between metals and non-metals
 1A elements are alkali metals
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– They are soft shiny metals that usually combine with
group VIIA nonmetals in chemical compounds in a
1:1 ratio.
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2A elements are the alkaline earth metals
– Both alkali and alkaline earth metals are chemically
reactive, but 2A metals are less reactive than 1As.
– They combine with the group VIIA nonmetals in a 1:2
ratio.
Transition Metals & Metalloids
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Transition metals
– The remaining 1-8B elements are all transition
elements
– The transition elements also have valence
electrons in two shells instead of one.
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Inner transition metals
– The lanthanide and actinide series comprise
the inner transition metals
Metalloids
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Metalloids have characteristics of both
metals and nonmetals and so can’t be
classified as either, but something in
between.
– They are good conductors of heat and
electricity
– They are not good conductors or insulators.
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The six metalloids are B, Si, Ge, As, Sb,
and Te.
Nonmetals
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There are 17 nonmetals in the periodic table,
and they are characterized by four major
physical properties.
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–
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They rarely have metallic luster.
They are usually gases at room temperature.
Nonmetallic solids are neither malleable nor ductile.
They are poor conductors of heat and electricity.
The elements above the B  At stairstep are
nonmetals
Nonmetals (cont)
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Group 6A contains the chalcogen elements
Group 7A contains the highly reactive halogen elements
– They are fluorine, chlorine, bromine, and iodine.
– The halogens exist as diatomic molecules in nature.
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Group 8A comprises the completely non-reactive noble
gases
– The noble gases are also called rare gas elements, and they all
occur in nature as gases.
– The noble gases fulfill the octet rule by having a full outer level
with 8 valence electrons.
– Therefore, they do not undergo chemical reactions because they
do not accept any electrons.
Valence Electrons and the Periodic
Table
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Valence Electrons and Group
– Atoms in the same group have the same chemical
properties because they have the same number of
valence electrons.
– Moreover, they have the same outermost orbital
structure
 E.g. 1A elements all have s1 valence electrons
 E.g. 2A elements all have s2 valence electrons
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Valence Electrons and Period
– The primary quantum number (n) for an element’s
valence electrons is the same its period.
 E.g. Lithium’s valence electron is n=2 and Li is found in the
2nd period
The Octet Rule
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Atoms tend to lose, gain, or share
electrons until they are surrounded by 8
valence electrons
Exceptions to Aufbau Order
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Subshell degeneracies occur in elements
larger than Vanadium
– i.e. different 4s and 3d orbitals have nearly
the same energy
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Also, it turns out that full and half-full
sublevels have the most stability.
Exceptions: Copper (Cu), Silver
(Ag), Gold (Au)
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Strict Aufbau ordering of Cu would be [Ar]4s23d9
– experimental observation shows this to be an excited
state
– the ground state has a configuration of [Ar]4s13d10
The observed configuration for Cu creates a ½full s and a full d, which is more stable than a
full s and a partial d
 Ag is NOT [Kr]5s24d9, but [Kr]5s14d10
 Au is NOT [Xe]6s25d9, but [Kr]6s15d10
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Exceptions: Lanthanum and
Actinium
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Aufbau would place them in the inner
transition series, but instead they are in
the scandium family
– i.e. La is [Xe]6s25d1
– i.e. Ac is [Rn]7s26d1
Exceptions:
Chromium (Cr), Molybdenum (Mo),
but NOT Tungsten (W)
Cr is [Ar]4s13d5, NOT [Ar]4s23d4
 Mo is [Kr]5s14d5, NOT [Kr]5s24d4
 W IS [Xe]6s15d5
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Ionization Energy
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The quantity of energy
required to remove an electron
from the gaseous atom or ion.
For Aluminum
Al (g)
kJ/mol
 Al + (g)
kJ/mol
 Al2+ (g)
kJ/mol
 Al3+ (g)
kJ/mol
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Al+ (g) + e-
I1 = 580
Al2+ (g) + e-
I2 = 1850
Al3+ (g) + e-
I3 = 2740
Al4+ (g) + e-
I4 = 11,600
Periodic Trends
First ionization energy:
 increases from left to right across
a period Why?
 decreases going down a group.
Why?
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Trends in Ionization Energies for
the Representative Elements
The Values of First Ionization Energy for
the Elements in the First Six Periods
Question
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The first ionization energy for
the group IIA elements are
significantly higher than those of
the Group IA elements in the
same periods. Why?
Question
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The first ionization energy of
the Group IIIA elements are lower
than the IIA elements in the same
period. Why?
Question
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Group VIA elements have
slightly lower first ionization
energies than Group VA elements
in the same period. Why?
Electron Affinity
The energy change
associated with the addition of
an electron to a gaseous atom.
 X(g) + e  X(g)
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Note: the more negative the
electron affinity, the more
energy is released
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The Electronic Affinity Values for Atoms
Among the First 20 Elements that Form
Stable, Isolated X- Ions
 Helium
Questions
and Beryllium do not form
stable isolated negative ions. Why?
 Nitrogen
does not form a stable,
isolated N- (g) ion, whereas carbon
forms C-(g). Why?
 In
contrast to nitrogen, oxygen can
add an electron to form the stable Oion. Why?
Periodic Trends
Atomic Radii:
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decrease going from left to
right across a period; Why?
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increase going down a group.
Why?
The Radius of an Atom
Atomic Radii for
Selected Atoms
Ionic Radii
 What
is the trend for ionic radii?
 Which of the Period 3 ions would
be the smallest?
 Na+, Mg2+, Al3+, S2-, Cl-
Sizes of Ions Related to Positions of
the Elements in the Periodic Table
Electronegativity Increases
Up and To the Right
What is electronegativity?
How tightly an atom holds on to its
valence electrons.
 Essentially, this value depends on
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– the number of positively charged protons in
the atom’s nucleus
– the radius of the outermost electron shell
What is electronegativity?
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The more positive the nucleus
The smaller the valence electron shell around it
The greater the attraction between nucleus and
electrons
Thus, the more electronegative the atom!
Thus, a high electronegativity value implies that the
valence electrons are tightly held and require a large
amount of energy to remove.
Oxidation Numbers
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The Octet Rule states that atoms want to have
their valence shell filled with electrons.
– This means that, ideally, atoms are most stable with 8
valence electrons
– N.B. This is not true for Period 1. Why?
Atoms will gain or lose electrons to form ions in
order to fulfill the Octet Rule.
 The charge they take on in this process is called
the valence.
 The oxidation state is, for ions, equal to the
valence.
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