Chemistry Ch. 5

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Chemistry Ch. 5
Periodic Law
5-1 History of the Periodic Table
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1860-Cannizzaro discovered a method for
accurately measuring Relative Masses,
which led chemists to agree on standard
values for atomic masses of elements
This led to a search for a relationship
between mass and other properties of the
then 60 discovered elements
Mendeleev
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Mendeleev wanted to organize the
elements according to their properties so
he wrote each element on a card, along
with its mass and its chemical and physical
properties and looked for trends/patterns.
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He noticed that when he put them in
order of increasing atomic mass, certain
similarities appeared in their chemical
properties at regular intervals (periodic).
1869- 1st periodic table pg. 124
Placed in vertical groups of increasing
atomic mass so that the horizontal rows
resembled each other chemically
Mendeleev’s First Periodic Table
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This table left empty places in the table
where he felt something should be
Mendeleev is given credit for the periodic
law (We’ll discuss periodic law in a bit)
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Two questions remained regarding
Mendeleev’s table:
Why could most of the elements be
arranged in the order of increasing
atomic mass but a few could not?
What was the reason for chemical
periodicity?
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1911-(40 years later and after discovering
all the subatomic particles) Moseley
noticed that the elements fit better with
patterns when he placed them in order of
increasing number of protons (atomic
number).
Periodic Law-The physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
Modern Periodic Table
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Chemists have discovered new elements
to fill the spaces and even synthesized
new ones
Periodic Table-an arrangement of the
elements in order of their atomic numbers
so that elements with similar properties
fall in the same column or group
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Noble Gases: Group 18-chemically unreactive
Lanthanides: 14 elements with atomic numbers
from 58 to 71 with similar physical and chemical
properties
Actinides: 14 elements with atomic numbers
from 90 to 103
These last two groups belong in periods 6-7 but
are set off below the main table to save space
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Periodicity can be seen in any group of the
periodic table
Fig. 5-4
The differences in each element of Group
1 and Group 18 are the same
The elements in Group 2 and Groups 1317 show a similar pattern.
5-2 Electron Configuration and the
Periodic Table
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Chemical properties are a result of the
electron configuration of an element
Ex: Noble Gases-the highest energy level
is occupied, making the noble gases stable
and unreactive
-highest configuration level governs the
properties of the element
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Groups-Vertical column of elements, show
related elements with similar properties
Periods-Horizontal row of elements, the
length depends on the sublevel being
filled
Principal Quantum number of the highest
occupied energy level tells us the period
the element is in
(# protons)
Atomic No.
14
Groups (down)
Si
Element Symbol
28.086
Atomic Mass
Silicon
Element Name
Atoms are Neutral:
(#protons = # electrons)
How many protons? 14
Periods (across)
How many electrons? 14
Sublevel Blocks
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Fig. 5-5 Based on the electron
configurations of the elements, the
periodic table can be divided into four
sublevel blocks
Pg. 129
Sublevel Periodic Table
Groups
Alkali metals-Group #1 Li-Fr
-silvery appearance and are soft enough to cut
with a knife
-so reactive, they are not found in nature and
have to be stored in kerosene because they
react violently with air or moisture.
-melting point decreases down the group
-have 1 electron in outer energy level
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Groups or Family Names
18) NOBLE GASES
Inner earth metals
17) HALOGENS
1) ALKALI METALS
2) ALKALINE METALS
TRANSITION
METALS
Alkaline-earth metals- Group #2 Be-Ra
-harder, denser, and stronger than alkali
metals
-have higher melting points than Group #1
-less reactive than Group #1, but still too
reactive to be found in nature
-have 2 electrons in outer energy level
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Hydrogen and Helium
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Hydrogen does not share the same
properties as Group #1; does not share
properties with any other group in the
table
Helium does not behave like Group #2 but
like Group #18 so it is part of the Noble
Gases (Group 18).
Transition Elements
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D-block elements that are metals with
typical metallic properties
-good conductors of electricity and a high
luster
-less reactive than alkali metals or alkaline
earth metals
-do not easily form compounds, exist as
free elements
Main Group Elements
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P-block elements together with the s-block
elements
-properties vary greatly
-number of electrons equals the group #10
Halogens
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Group 17: F-At
-most reactive nonmetals
-react vigorously with metals to form salts
7 electrons in the outer energy level
Lanthanides and Actinides
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F-block elements
Lanthanides-shiny metals similar reactivity
to Group 2 elements
Actinides-all radioactive
Metals
Properties of metals and nonmetals
Metals
Luster (shiny silvergray, gold, copper)
Nonmetals
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Dull (all different
colors)
Malleable (hammered) Brittle
Ductile (drawn into
thin wire)
Not ductile
High Density & High
Melting Points
Low Density & Low
Melting Points
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Conductors of heat & Poor Conductors of
electricity
heat & electricity
Lose electrons
Gain electrons
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Metalloids
have both
metallic and
nonmetallic
characteristics
–along the
staircase.
Boron, Silicon,
Germanium,
Arsenic,
Antimony,
Tellurium,
Astatine,
Used in
electronics
5-3 Electron Configuration and
Periodic Properties
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Atomic Radii-as one-half the distance
between the nuclei of identical atoms that
are bonded together
Period-smaller atomic radii across a period
due to increasing the positive charge of
the nucleus
Group-increase down a group due to
increase in size of the atom
Ionization Energy
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Ion-atom(s) that has a positive or
negative charge due to gain/loss of an
electron
Ionization-process that results in the
formation of an ion
Ionization energy-the energy required to
remove one electron from a neutral atom
of an element
Ionization Energy
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Period-increase across a period, generally
Group-decrease down the group, generally
Table 5-3
Highlighted energy reflects the energy
needed to remove an electron from the
noble gas configuration or the full energy
level
Electron Affinity
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The energy change that occurs when an electron
is acquired by a neutral atom
Most atoms release energy when they acquire
an electron
Energy released represented by a negative
number ex: Group 17
Energy absorbed represented by a positive
number ex: Group 1 or Group 18
This book uses negative numbers to represent
the energy released
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Period-generally, become more negative
across each period (increase-release more
energy as we go across because we
acquire more electrons )
Group-not as regular, electrons add with
greater difficulty down a group (decreaserelease less energy going down because
acquire fewer electrons)
Ionic Radii
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Cation-positive ion
Anion-negative ion
Period-metals at the left tend to form
cations and the nonmetals tend to form
anions
Cationic radii decrease across a period due
to shrinking electron cloud
Anionic radii decrease across a period
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Group-gradual increase of ionic radii down
a group
Valence Electrons
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The electrons available to be lost, gained,
or shared in the formation of chemical
compounds (in the outer shell)
Groups 1 & 2 have 1 & 2 electrons,
respectively
Groups 13-18 have valence electrons
equal to group # - 10
Electronegativity
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A measure of the ability of an atom in a
chemical compound to attract electrons
Period-increase across each period, with
some exceptions
Group-decrease down a group or stay
about the same
Highest electronegative element=Fluorine
Electronegativity numbers
Trends

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