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Chemistry Chapter 5
THE PERIODIC LAW
Mendeleev’s Periodic Table
Dmitri Mendeleev
Mendeleev – organized periodic table
 Tried to organize periodic table according to
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properties
Vertical columns in atomic mass order
Made some exceptions to place elements in rows
with similar properties (Tellurium and Iodine)
Horizontal rows have similar chemical properties
Gaps for “yet to be discovered” elements
Left questions: why didn’t some elements fit in
order of increasing mass? Why did some elements
exhibit periodic behavior?
Henry Moseley
 Discovered that periodic table was in atomic number
order, not atomic mass order
 Explained the Te-I anomaly
 Determined the atomic numbers using x-rays
Periodic Law
 Physical and chemical properties of the elements are
periodic functions of their atomic numbers
Modern Periodic Table
 Discovery of noble gases yields new family (Group 18
– aka inert gases)
 Lanthanides (#58 - #71)
 Actinides (#90 – #103)
Periods and Blocks of the Periodic Table
 Periods – horizontal rows
 Corresponds to highest principal quantum number
 Groups/Families – vertical columns; these elements
share similar chemical properties (they have the
same number of valence electrons)
 Blocks – periodic table can be broken into blocks
corresponding to s, p, d, f sublevels
Orbital filling table
s block
 S block includes groups 1 and 2
 Group1 – “The alkali metals”
 One s electron in outer shell
 Soft, silvery metals of low density and low melting
points
 Highly reactive, never found pure in nature
s block
 Group 2 – “Alkaline Earth Metals”
 2 s electrons in outer shell
 Denser, harder, stronger, less reactive than Group 1
 Too reactive to be found pure in nature
Periodic Table with Group Names
The Properties of a Group:
the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
React with halogens to form salts
d block
 Groups 3 -12
 Metals with typical metallic properties
 Referred to as transition metals
 Group number = sum of outermost s and d electrons
p block
 Groups 13-18
 Properties vary greatly – metals, metalloids, and
nonmetals
 Group 17 – halogens are most reactive of non metals
 Group 18 – noble gases are NOT reactive
f block
 Lanthanides – shiny metals similar to group 2
 Actinides – all are radioactive; plutonium –
lawrencium are man-made
Properties of Metals
 Metals:
 good conductors of
heat and electricity
Malleable
Ductile
 Have high tensile
strength
 luster
Examples of Metals
Potassium, K
reacts with
water and
must be
stored in
kerosene
Copper, Cu, is a
relatively soft metal,
and a very good
electrical conductor.
Mercury, Hg, is the only
metal that exists as a
liquid at room
temperature
Zinc, Zn, is
more stable
than potassium
Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a
great example of a nonmetallic element.
 Nonmetals are:
 poor conductors of heat and electricity
 brittle
Many are gases at room temperature
Examples of Nonmetals
Sulfur, S, was
once known as
“brimstone”
Graphite is not the only
pure form of carbon, C.
Diamond is also carbon;
the color comes from
impurities caught within
the crystal structure
Microspheres
of
phosphorus,
P, a reactive
nonmetal
Properties of Metalloids
Metalloids straddle the
border between metals
and nonmetals on the
periodic table.
 They have properties of both metals and
nonmetals.
Metalloids are more brittle than metals, less
brittle than most nonmetallic solids
 Metalloids are semiconductors of electricity
Many used in computer parts
 Some metalloids possess metallic luster
Silicon, Si – A Metalloid
 Silicon has metallic luster
 Silicon is brittle like a
nonmetal
 Silicon is a semiconductor of
electricity
Other metalloids include:
 Boron, B
 Germanium, Ge
 Arsenic, As
 Antimony, Sb
 Tellurium, Te
IONS
 Cation- positively charged ion
 Anion- negatively charged ion
IONS
 Group 1
 1 valence electron (+1)
 Will lose 1 electron
 Group 2
 2 valence electrons (+2)
 Will lose 2 electrons
 Group 13
 3 valence electrons (+3)
 Will lose 3 electrons
D block IONS
 Cannot predict
 Memorize:
 Ag+
 Cd 2+
 Zn 2+
IONS
 Group 14
4 valence e- (4+ or 4-)
 Will lose 4 or gain 4 electrons
Group 15
 5 valence e Gains 3 electrons (3-)
Group 16
 6 valence e Gains 2 electrons (2-)
Group 17
 7 valence e Gains 1 electron (1-)
Group 18
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
No ions
Determination of Atomic Radius:
Atomic Radius: Half the diameter of an
atom
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to increased number of protons and electrons
Radius increases down a group
Addition of principal quantum levels
Table of Atomic Radii
Ionization Energy - the energy required to remove an
electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Increased number of protons and
electrons leads to greater attraction
Tends to decrease down a group
Outer electrons are farther from the
nucleus
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Another Way to Look at Ionization Energy
Ionic Radii
Cations
Anions
Positively charged ions
Smaller than the corresponding
atom
Negatively charged ions
Larger than the corresponding
atom
Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase
across a period
* more positive charges in the nucleus
to attract more elctrons
Electronegativities tend to decrease
down a group or remain the same
* additional energy levels result in
less attraction to the nucleus