PERIODIC TABLE
The most awesome chemistry tool ever!
UNIT OBJECTIVES:



Understand the history of who built up the
periodic table, how they did it, and what law was
made
Become familiar with structure of periodic table,
how the e- is related to the structure of the table,
and properties of chemicals based on their
location
Discover trends related to electron configuration
and periodic properties; use electron
configuration to predict location on periodic table.
HISTORY
•
Explain the roles of Mendeleev and Moseley in the
development of the periodic table.
•
Describe the modern periodic table.
•
Explain how the periodic law can be used to
predict the physical and chemical properties of
elements.
•
Describe how the elements belonging to a group
of the periodic table are interrelated in terms of
atomic number.
MENDELEEV AND PERIODICITY
•
In the late 1800s, scientist Mendeleev noticed:
•
when the elements were arranged in order of
increasing atomic mass, certain similarities in their
chemical properties appeared at regular intervals and
formed a pattern.
•
•
Repeating patterns are referred to as periodic.
Something that occurs periodically occurs at
regular/fixed intervals
MENDELEEV AND PERIODICITY
•
Not all elements had been discovered, so when
arranging the table by atomic mass,
•
He left spaces for future elements to be discovered
and placed into the table based on their mass
•
He even predicted what their properties would
likely be
MENDELEEV AND PERIODICITY

Where Mendeleev’s work left off there were some
questions:

Why could most elements be arranged in the order of
increasing atomic mass and other not?

What was the reason behind chemical periodicity?
MOSELEY AND THE PERIODIC LAW


Moseley continued research on the periodic table
and was able to propose an answer to the first
question.
Elements follow a clearer pattern when arranged
by their nuclear charge (# of protons) instead of
their nuclear mass!
Why do you think this is?
 Hint: remember Hydrogen-3 and Helium-3 and their
similarities and differences?

MOSELEY AND THE PERIODIC LAW

A big example of what Moseley did: he realized
that Tellurium (which has an atomic mass of
about 127.6 amu) should be before Iodine
(which has an atomic mass of 126.9 amu) on the
periodic table because Tellurium has the atomic
number of 52 and Iodine has the atomic number
of 53.
MOSELEY AND THE PERIODIC LAW



Moseley led to the editing of Mendeleev’s
principle of chemical periodicity:
Periodic law: the physical and chemical
properties of the elements are periodic functions
of their atomic numbers.
(When elements are arranged by increasing
atomic number, similar properties occur in
elements at regular intervals)
MODERN PERIODIC TABLE


After all the elements had been discovered (some
synthesized as well) our modern periodic table
was been established
Periodic table: an arrangement of the elements in
order of their atomic numbers so that elements
with similar properties fall into the same column
or “group”.
SPECIFIC GROUP: NOBLE GASES

Discovered in the late 1800s, the noble-gases
were slowly discovered, often due to indirect
observation

Helium was discovered due to its emission
spectrum in sunlight

Argon was not-reactive and was deduced when
mass of an unknown gas could not be accounted for
when all other gases had been reacted out.
LANTHANIDES AND ACTINIDES HAVE
THEIR OWN BLOCK

Lanthanides:
The elements between Cerium (Atomic # 58) and
Lutenium (Atomic # 71) did not follow the properties
of the elements around them
 Whereas Lanthanum (atomic # 57) and Hafnium
(atomic # 72) do fit in their groups


Actinides:
The elements between atomic #s: 90 and 103
 Many are synthetic and they do not follow the
properties of the metals around them in the periodic
table

PERIODICITY
 We
will begin to explore the similar
properties groups of the periodic table
exhibit.
 Groups



to begin thinking about:
Alkali Metals (s-block)
Alkaline Earth Metals (s-block)
Main-group elements (p-block
Halogens (p-block)
 Noble Gases (p-block)


Transition metals (d-block)
ELECTRON CONFIGURATION AND
PERIODIC PROPERTIES
Objectives:
• Explain the relationship between electrons in
sublevels (shapes) and period length in the periodic
table.
•
Locate and name the four blocks of the periodic table.
Explain the reasons for these names.
•
Discuss the relationship between group configurations
and group numbers.
•
Describe the locations in the periodic table and the
general properties of the alkali metals, the alkalineearth metals, the halogens, and the noble gases.
PERIODS AND BLOCKS OF THE PERIODIC
TABLE

Periods are the horizontal rows of the periodic
table.
Elements are arranged vertically in the periodic
table in groups that share similar chemical
properties.
 The four blocks: the s, p, d, and f blocks


each block corresponds to the electron sublevel being
filled in that block.
PERIODS & BLOCKS OF THE
PERIODIC TABLE
Period #
Sublevels
# of elements in
(shapes) in order period
of filling
1
1s
2
2
2s 2p
8
3
3s 3p
8
4
4s 3d 4p
18
5
5s 4d 5p
18
6
6s 4f 5d 6p
32
7
7s 5f 6d 7p
32
S-BLOCK ELEMENTS
•
•
The elements of Group 1 of the periodic table are
known as the alkali metals.
•
lithium, sodium, potassium, rubidium, cesium, and
francium
•
In their pure state, all of the alkali metals have a
silvery appearance and are soft enough to cut with a
knife.
The elements of Group 2 of the periodic table are
called the alkaline-earth metals.
•
•
beryllium, magnesium, calcium, strontium, barium, and
radium
Group 2 metals are less reactive than the alkali metals,
but are still too reactive to be found in nature in pure
form.
HYDROGEN AND HELIUM

Hydrogen is NOT a metal;


Its properties are very different from the metals in
the alkali metal group
Helium fits in with the Noble Gases as it has a
full valence shell and is not at reactive

Alkaline earth metals have a full s-sublevel, but their
whole valence shell is not full (their p-sublevel is
empty)

A valence shell is the highest energy level that
electrons occupy. Elements are more stable/less
reactive when their valence shells are full of
electrons
THE d-BLOCK ELEMENTS (GROUPS 3-12)

Transition metals: the d-block elements are
metals with typical metallic properties
Ductility
 Malleability
 Conductivity
 Having luster


Groups and periods within the transition metals
have varied properties and it is therefore hard to
group metals by individual properties
p-BLOCK ELEMENTS (GROUP 13-18)



Main-group elements: the elements in the p-block
of varying properties
All non-metals are in the p-block (aside from
hydrogen)
All metalloids (boron, silicon, germanium,
arsenic, antimony, and tellurium) are in the pblock
RELATIONSHIPS BETWEEN GROUP #,
BLOCKS, AND ELECTRON CONFIG.
Group # Group
Block
configuration
1,2
3-12
13-18
ns1, ns2
(n - 1)d1-10ns0-2
ns2np1-6
s
Comments
1 or 2 electrons in ns sublevel
d
The sum of electrons in ns and (n-1)d
equals group number
p
Number of electrons in np sublevel
equals group number minus 12
HALOGENS

Elements in group 17 are the halogens


They are highly reactive in similar ways


Fluorine, chlorine, bromine, iodine, and astatine
All form salts with group 1 and 2 metals
They are volatile ( gas forms easily)
METALLOIDS

Metalloids have properties of both metals and
non-metals



boron, silicon, germanium, arsenic, antimony, and
tellurium
Semiconducting, brittle, harder than most
metals, have an opalescent luster
Relatively reactive (Bi only found in elemental
form)
f-BLOCK ELEMENTS
LANTHANIDES AND ACTINIDES



Lanthanides are shiny metals and have similar
reactivity to group 2 metals
Their position reflects the fact that they involve
the filling of the 4f sublevel
Actinides:

Only 4 are found in nature (thorium through
neptunium)
Alkali Metals
Alkaline Earth Metals
Other Metals
Metalloids
Inner transition
Metals
Other Non-Metals
Halogens
Noble-Gases
DO NOW:
•
List properties of the
following elemental groups:
•
•
•
•
•
•
•
Table 1: Halogens
Table 2: Noble Gases
Table 3: Alkali Metals
Table 4: Alkaline Earth
Metals
Table 5: Transition metals
Table 6: Actinides
Describe what you know
about an element by looking
at its position in the
periodic table.
•
•
•
•
•
•
•
Share with your table
•
Table 1: Br
Table 2: Ar
Table 3: K
Table 4: Ca
Table 5: Cu
Table 6: Fe
Identify any noticeable
trends.
DO NOW
If you were absent Friday:




How did Mendeleev organize
the periodic table?
If you were here Friday:

Why didn’t this always work?
How did Moseley organize the
periodic table?
What are the groups of the
periodic table?

Include at least 1 property for
each group


Please take out your practice
problem sheet and complete
the problems for Atomic
radius and ionization energy
using your graphic organizer
Look over your SRQ Answers
Help your neighbor with their
do now if they were absent
Friday
ATOMIC RADII
(THE DISTANCE BETWEEN THE NUCLEUS
& THE END OF THE CLOUD)

One half the distance between the nuclei of
identical atoms that are bonded together
Atomic radii: 200 pm
Bond
length
400 pm
ATOMIC RADII
Increases in size from right to left
 Increases in size from up to down

IONIZATION ENERGY
(HOW TO KNOCK AN ELECTRON OFF!)



An ion is an atom that

has more or less electrons than it has protons.

is either positively (cation) or negatively (anion)
charged
Ionization is anything that gives matter a
charge
Ionization Energy (IE) is the amount of energy
needed to remove an electron from an atom
IONIZATION ENERGY CONTINUED

Ionization often occurs because
The valence electrons are more stable in the ionic
configuration
 To complete a valence shell


Gaining an electron can cause a neutral atom to
lose or gain energy depending on the atom
A metal gaining an electron would absorb energy and
become less stable
 A non-metal gaining an electron would release
energy and become more stable

IONIZATION ENERGY
(HOW TO KNOCK AN ELECTRON OFF!)

Ionization energy trends:
Increases from left to right
 Increases from down to up

ADDITIONAL IONIZATION ENERGIES


Removing a second electron is the 2nd Ionization
Energy, and it typically greater than the first
ionization energy
With each additional electron removed, the
Ionization Energy will increase.
ELECTRON AFFINITY

The change in energy associated with the addition of an
electron to a neutral atom

Green means no electron affinity

Yellow means high electron affinity (trends by groups)
IONIC RADII
Positive ion = cation
 Negative ion = anion

IONIC RADII
VALENCE ELECTRONS

Valence electrons are available to be lost, gained,
or shared during chemical reactions and
compound formation.
Group # Group e- Config
# of valence e-
1
ns1
1
2
ns2
2
13
ns2p1
3
14
ns2p2
4
15
ns2p3
5
16
ns2p4
6
17
ns2p5
7
18
ns2p6
8
ELECTRONEGATIVITY

Electronegativity how much an atom in a
compound attracts electrons from another atom
in the compound

(how much the atom is greedy for electrons)
PERIODIC PROPERTIES OF THE
d-BLOCK AND f-BLOCK

Atomic radii decrease from left to right

Ionization energy increases from left to right

Electronegativity slightly increase from left to
right
-1
-2
-3
+3
+2
+1
COMMON ELEMENTAL IONS
8.2
WHY?


A lot of these trends are related to electron
shielding
Electron shielding is how the inner electrons
effect how the outer electrons act due to charge
different and repulsion