Chapter 5
Periodic Law
History of the Periodic Table
SECTION 1
Mendeleev and Chemical
Periodicity
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When the Russian chemist
Dmitri Mendeleev heard
about the new atomic
masses he decided to
include the new values in a
chemistry textbook that he
was writing
Mendeleev hoped to
organize the elements
according to their
properties
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He placed the name of each known element
on a card, together with the atomic mass of
the element and a list of its observed physical
and chemical properties
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He then arranged the cards according to
various properties and looked for trends or
patterns
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Mendeleev noticed that when the elements
were arranged in order of increasing atomic
mass, certain similarities in their chemical
properties appeared at regular intervals
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Such a repeating pattern is referred to as
periodic
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Mendeleev’s procedure left several empty
spaces in his periodic table
In 1871, he predicted the existence and
properties of the elements that would fill three
of the spaces
By 1886, all three elements had been
discovered
scandium, Sc, gallium, Ga, and germanium, Ge
Their properties are very similar to those
predicted by Mendeleev
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Success of Mendeleev’s predictions persuaded
most chemists to accept his periodic table and
earned him credit as the discoverer of the
periodic law
Two questions remained
(1) Why could most of the elements be arranged
in the order of increasing atomic mass but a few
could not?
(2) What was the reason for chemical
periodicity?
Moseley and the Periodic Law
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In 1911  English
scientist Henry
Moseley examined the
spectra of 38 different
metals
Moseley discovered a
previously unknown
pattern
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The elements in the periodic table fit into
patterns better when they were arranged in
increasing order according to nuclear charge,
or the number of protons in the nucleus
Moseley’s work led to both the modern
definition of atomic number and the
recognition that atomic number, not
atomic mass, is the basis for the
organization of the periodic table
Electron Configuration and the Periodic
Table
SECTION 2
The Modern Periodic Table
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“Periodic” - Repeating patterns
Listed in order of increasing number of protons (atomic
#)
Properties of elements repeat
Periodic Law- “the physical and chemical properties of
the elements are periodic functions of their atomic
numbers.”
Periods and Blocks of the
Periodic Table
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Elements are arranged vertically in the
periodic table in groups that share similar
chemical properties
They are also organized horizontally in
rows, or periods
The length of each period is determined by
the number of electrons that can occupy the
sublevels being filled in that period
main group elements
Metals
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Most solids (Hg is liquid)
Luster – shiny.
Ductile – drawn into thin wires.
Malleable – hammered into
sheets.
Conductors of heat and
electricity.
Include transition metals –
“bridge” between elements on
left & right of table
Non-Metals
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Properties are generally opposite of
metals
Poor conductors of heat and
electricity
Low boiling points
Many are gases at room
temperature
Solid, non-metals are brittle (break
easily)
Chemical properties vary
Metalloids
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stair-step pattern
Have properties similar
to metals and non-metals
Ability to conduct heat
and electricity varies with
temp
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Better than non-metals but
not metals
semiconductors
Group 1 – The Alkali Metals
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The elements of Group 1
of the periodic table
(lithium, sodium,
potassium, rubidium,
cesium, and francium)
are known as the alkali
metals
In their pure state, all of
the alkali metals have a
silvery appearance and
are soft enough to cut
with a knife
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Because they are so reactive, alkali
metals are not found in nature as free
elements
They combine strongly with most
nonmetals
And they react strongly with water to
produce hydrogen gas and aqueous
solutions of substances known as
alkalis
Because of their extreme reactivity with
air or moisture, alkali metals are
usually stored in kerosene or oil
Group 2 – The Alkaline-Earth Metals
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The elements of
Group 2 of the
periodic table are
called the alkalineearth metals
Atoms of alkalineearth metals
contain a pair of
electrons in their
outermost s
sublevel
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Group 2 metals are harder, denser, and
stronger than the alkali metals
They also have higher melting points
Less reactive than the alkali metals, but also
too reactive to be found in nature as free
elements
Transition Elements
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Good conductors of electricity and have a
high luster
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They are typically less reactive than the
alkali metals and the alkaline-earth metals
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Some are so unreactive that they do not
easily form compounds, existing in nature
as free elements
Mercury
Tungsten
Vanadium
uranium
Rare Earth Elements
Lanthanide series (period 6)
 Actinide Series (period 7)
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Some radioactive
Separated from table to make easy to
read/print
silver, silvery-white, or gray metals.
Conduct electricity
Halogen Family (“salt-former”)
-7 Valence Electrons
-most active nonmetals
-never found pure in nature
-react with alkali metals easily
(forms salts)
-F most active halogen
Halogens cont…
F
compounds in toothpaste
 Cl kills bacteria
 I keeps thyroid gland working
properly
bromine
The Noble Gases (Inert Gases)
- non-reactive
- outermost e- shell is
full (8 VE)
- In “neon” lights
-in earth’s
atmosphere (less
than 1%)
Neon
Section 5.3
Electron Configuration and
Periodic Properties
Periodic Trends
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In periodic table, there is a DECREASE in
atomic radii across the periods from left to
right
Caused by increasing positive charge of
nucleus (more protons = more positive
charge)
Group Trends
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Radii of elements decrease as you go UP a
group
Electrons occupy sublevels in consecutively
higher main energy levels (located further
away from nucleus)
In general, the atomic radii of the maingroups elements decrease from the
bottom to the top of a group
2. Ionization Energy
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Electrons can be removed from an atom if
enough energy is supplied
Using A as a symbol for an atom of ANY
element, the process can be expressed as
follows:
A + energy  A+ + e-
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A+ represents an ion of element A with single
positive charge (a 1+ ion)
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Ion  an atom or group of bonded atoms
that have a positive or negative charge
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Ionization  any process that results in the
formation of an ion
Period Trends
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In general,
ionization energies
of the main-group
elements INCREASE
across each period
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Caused by increasing
nuclear charge
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Higher charge more
strongly attracts
electrons in same
energy level
Group Trends
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Ionization energies generally INCREASE
going UP a group
Electrons going down in group are in higher
energy levels, so further away from the
nucleus
Removed more easily
Also more electrons between outermost
electrons and the nucleus (shields them from
attraction to positive nucleus)
What are Valence electrons?
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outermost e-’s
Responsible for chem props
Elements in same group…
same # of VE
ALL atoms want full outer
energy level (usually 8 VE)
To get full outer energy level, some
elements:
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lose e- (metals)
gain e- (non-metals)
share electrons (some non-metals & metalloids)
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Main-group elements – valence e- are in
outermost s and p sublevels
Inner e- held too tightly by nucleus to be involved
in compound formation
6. Electronegativity
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Valence e- hold atoms together in
compounds
In many compounds, negative charge
centered around one atom more than another
Uneven distribution of charge has effect on
compound’s properties
Useful to have measurement of how strongly
one atom attracts e- of another atom
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Electronegativity  measure of the ability of
an atom in a chemical compound to attract
electrons
Most e-neg element (fluorine) – randomly
assigned value of 4
Other values calculated in relation to F
Period Trends
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e-negs tend to INCREASE across each
period
There are exceptions (of course)
Alkali and alkaline-earth metals are least eneg
In compounds, their atoms have low
attraction for e-
Group Trends
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Electronegativities tend to INCREASE
going UP a group or stay the same
At higher energy levels electrons being
added are further away from the nucleus
Therefore, less attraction to the nucleus
Also more electrons between outermost
electrons and the nucleus (shields them from
attraction to positive nucleus)