The Periodic Table
History of the Table
Periodic Law
Periodic Trends
Development of the Periodic
Table
• Dmitri Mendeleev organized the first
periodic table according to atomic masses.
• He noticed that chemical properties
repeated at regular intervals, thus the term
periodic.
• A few elements contradicted the pattern, so
he rearranged them so their properties were
aligned, even though the atomic mass was
not in increasing order.
Predicting the Future?
• The most amazing thing about Mendeleev’s
work is that he predicted the properties of
elements that did not exist yet.
• He left empty spaces for these elements,
which were found and named Scandium,
Gallium and Germanium. The properties
that he predicted were very similar to the
actual properties observed.
Moseley and the Periodic Law
• Though most of the
elements could be
arranged according to
atomic mass, not all fit
this pattern.
• Moseley worked with the
spectra of elements and
determined that they fit
better into the pattern
when arranged according
to atomic number.
• Moseley’s periodic table
still adhered to the
chemical periodicity that
Mendeleev observed.
• Therefore it is Mendeleev
who is credited with
periodic law, which states
that the physical and
chemical properties of
elements are periodic
functions of their atomic
numbers.
Additions to the Table
• More than 40 new elements have been
added since Moseley’s time.
• Most significant additions:
• Noble gases - Group 18
• Lanthanides (#58-71)
• Actinides (#90-103)
Organization of the periodic
table
• Rows are referred to as periods.
• Columns are referred to as groups or
families.
• We have discussed the organization of the
periodic table as a result of electron
configuration.
• There are also patterns of periodicity
between elements in a group.
Organization of the periodic
table
• Valence electrons play an important role in
reactivity of elements. Valence electrons
are defined as the electrons in the outermost
energy level.
• You will notice a pattern between the
groups on the periodic table and number of
valence electrons.
The Octet Rule
• Elements want to have 8 electrons in their
valence shell.
• To get 8 electrons, some atoms gain
electrons and others lose electrons.
• When this happens, an ion is formed.
• Ions are atoms that have gained or lost one
or more electrons.
• The only exceptions to the octet rule are
Hydrogen and Helium…they desire 2
electrons to fill their valence shell.
Main group elements
• The term main group elements refers to
elements in the s and p blocks only.
• Remember the trend that allows you to
determine charge of an ion for elements in
these groups:
• Group 1=+1 Group 14=+/- 4
• Group 2=+2 Group 15= -3 Group 17=-1
• Group 13=+3 Group 16=-2
Group 1- Alkali Metals
• This group of metals is extremely reactive! This is
due to the fact that they all have only 1 valence
electron.
• In their pure state, all members of this group are
silvery in appearance and soft enough to be cut
with a butter knife.
• Because they are so reactive, they are not found in
nature as “free” elements.
• They react readily with water, even moisture in
the air.
• Therefore, they must be stored in kerosene or
mineral oil.
Group 2- Alkaline Earth Metals
• These metals are harder, denser and
stronger than alkali metals.
• They have 2 valence electrons.
• They have higher melting points and are
less reactive than group 1 metals.
• However, they are still too reactive to be
found “free” in nature.
Special Cases: Hydrogen and
Helium
• Even though H has 1 valence electron, it
does not share the same properties as the
other group 1 elements.
• Can you think of some differences?
• Helium has 2 valence electrons, but you
will notice that it is not placed above group
2.
• It is placed over group 18 because it has a
full valence shell.
Transition Elements
• These are elements belonging to groups
3-12.
• The number of valence electrons varies for
these elements.
• These elements have metallic properties that
you might suspect: good conductivity of
electricity, high luster and good
conductivity of heat.
• They are also less reactive than Group 1 and
2 metals, so they are found in nature in free
form.
P-block Elements
• Groups 13-18, excluding Helium
• The p-block and s-block elements are
collectively called the main group
elements.
• For the p-block elements, the number of
valence electrons is equal to the group
number minus 10.
• The groups vary drastically in their
properties.
The P-block Divide
• The p-block elements are divided into three
sections by the stair-step line located in the
center of the block.
• Elements that border the line are called
metalloids or semiconductors. These
include boron, silicon, germanium, arsenic,
antimony and tellurium.
• To the left of those elements are metals.
• To the right of the metalloids are
nonmetals.
The Halogens
• Halogens are elements that belong to group
17 on the periodic table.
• They are the most reactive nonmetals, so
they react vigorously with metals to form
salts.
• Their reactivity is based on the fact that
they have 7 valence electrons. Since all
elements desire a full valence shell
(halogens desire to get one more!)
The Halogens
• Fluorine and chlorine
are gases at room
temp.
• Bromine is a reddish
liquid and can burn
your skin!
• Iodine is a dark purple
solid.
• Astatine is rare but is
known to be a solid.
Only small quantities
exist.
Justifying Periodic Trends
(insert term here…
• It is REALLY
electronegative, for
important that you
example) because it is
understand how to
justify a trend observed further to the right on
on the periodic table.
the periodic table.
• This is a really major •You MUST explain in
topic in AP Chem.
terms of nuclear charge, or
• AP will not accept that more energy levels. I have
an element is more
given you the proper
justification for each trend
on the slides to follow!
Periodic Trends
• Atomic Radius- used
to determine the size
of an atom
• Defined as 1/2 the
distance between the
nuclei of two identical
atoms that are bonded
together
• The trend for atomic
radius on the periodic
table is:
• Size of atoms decrease
across a period (as a result
of increasing positive
charge of nucleus, which
attracts electrons more
closely)
• Size of atoms generally
increase down a group
(electron energy levels are
further from the nucleus
and increase in number)
Periodic Trends
• Ionization energythe energy required to
remove an electron
from an atom
• Any time an electron
is removed from an
atom, it is referred to
as ionization.
• The trend for ionization
energy on the periodic
table is:
• IE of main group elements
increases across a period
(this is due to increased
nuclear charge).
• IE of main group elements
decreases down a group
(electrons are further from
the nucleus requiring less
energy to remove an
electron.
Shielding
• Part of the reason that large atoms can give
up valence electrons so easily to form ions
is an effect called shielding.
• Shielding refers to the fact that the inner
shell electrons shield the positive nuclear
charge from the outer electrons, making
them easier to remove.
• An atom can have more than 1 IE if more
than one electron is removed!!!
Periodic Trends
• Electron Affinity- the
energy change that
occurs when a neutral
atom acquires an
electron
• Trends on the periodic
table:
• Generally, EA
increases across the p
block period.
• An exception to this is
between groups 14
and 15. It decreases
slightly.
• Hard to establish
group trends for EA.
Periodic Trends
• Ionic radii- the radius • Trends on the periodic
of an atom once it has
table:
become an ion
• Generally, metals, to
• Positive ions are called
the left of the staircations, which are
step line, form cations.
formed from the loss
• Cationic radii decrease
of electrons.
across a period, due to
• Negative ions are
increased nuclear
called anions, which
charge.
are formed from the
gain of electrons.
Ionic Radii Continued
• Nonmetals, to the right
of the stair-step line,
form anions.
• Generally, ionic radii
of anions decreases
across a period.
• For groups on the
periodic table: ionic
radii increase down a
group, due to adding
more energy levels.
Periodic Trends
• Electronegativitymeasure of the ability
of an atom in a
chemical compound
to attract electrons
• The reason this is
necessary is that the
negative charge of
electrons in a
compound is not
evenly distributed
between the atoms.
• Most electronegative atom
is Fluorine, and has an
electronegativity value of
4 assigned to it. Values for
the other elements are
assigned in relation to F.
Electronegativity Continued
• Trends on the periodic table:
• Electronegativities tend to increase across a
period.
• They tend to decrease or remain the same down a
group.
• Noble gases are unique because many are not
assigned electronegativity values, since they are
relatively unreactive.
• When a noble gas does form a compound, the
electronegativity value is quite high, like F.
Periodic Trends in the d and f
block
• Atomic radii- decrease across a period in the d
block
• Ionization energies- generally decrease across the
period for both f and d, BUT d block elements
have IE that increase down each group
• Ionic radii and ion formation- order in which
electrons are removed is exactly the reverse order
from their electron configuration
• Electronegativity- d block ranges from 1.1-2.54
• F block range from 1.1-1.5