The Periodic Table and Periodic Law
CHAPTER 6
Development of the Modern Periodic Table
SECTION 6.1
PERIODIC TABLE DEVELOPMENT
Until the 1790s, only 23 different elements
were known, such as silver, gold, carbon,
oxygen, etc
 These elements had been known for hundreds
or thousands of years
 The 1800s brought forth many changes to the
scientific community, including an explosion in
the number of elements known
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By 1870, there were about 70 known elements
 With the increase in the number of elements,
came an increased need to organize them
 The man given the most credit for organizing
the elements is Dmitri Mendeleev
 He organized the elements in order of
increasing mass, and when he did this, he
noticed a pattern in the properties of the
elements
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Mendeleev arranged his periodic table similar to a
winning configuration of solitaire
 Within rows, the elements were arranged by
increasing mass
 Within columns, elements were arranged by
similar chemical properties
 The usefulness of Mendeleev’s table was
confirmed by the discovery of new elements that
matched predicted properties according to their
location on the periodic table
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MODERN PERIODIC TABLE
Mendeleev’s table was a good step forward,
but it was not completely correct
 It soon became apparent, based on certain
chemical properties, that some elements were
not in the correct order
 It was not until 1913 when Henry Moseley
discovered that each element has a unique
positive charge that this was corrected
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Mendeleev did not know about subatomic
particles when he made his periodic table
 The fact that each element has a unique
charge, while the mass can vary because of
isotopes caused the periodic table to be
rearranged
 The modern periodic table would then now be
arranged by the atomic number
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PERIODIC LAW
The modern periodic table is arranged by
atomic number
 When the periodic table is arranged in this
fashion, there is a periodic repetition of
chemical properties from row to row. This is
called periodic law

The rows on the periodic table are called
periods
 The columns on the periodic table are called
groups or families
 Groups can be numbered in two separate
methods – groups can be designated with a
number and a letter, or just numbered 1-18
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The A groups (IA-VIIIA) are called representative
elements because they represent a wide variety
of chemical and physical properties
 The B groups (IB-VIIIB) are the transition
elements. They transition from very metallic to
less metallic

Starting below boron, draw imagine a staircase
down to the bottom of the periodic table
 Elements below and to the left of this staircase,
but not touching (aluminum is the exception) are
metals
 Metals are generally shiny, solid at room
temperature, good conductors of heat and
electricity, malleable, and ductile
 The far left group (IA) is called the alkali metals
 The next group (IIA) is called the alkaline earth
metals
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Alkali metals and alkaline earth metals tend to
be chemically active
 The alkali metals are more reactive than
alkaline earth metals
 Within the group of alkali and alkaline earth
metals, reactivity increases as we go down the
group
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The B elements (transition elements) are divided
into two groups: the transition metals and the
inner transition metals
 The inner transition metals are the two rows
located below the main body of the periodic table
 Inner transition metals (specifically the lanthanide
series) are used as phosphors, substances that
emit light when struck by electrons
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To the right and above (but not touching) the staircase are
the nonmetals
Nonmetals are generally gases at room temperature, brittle,
dull, poor conductors of heat and electricity, and nonductile
The group VIIA nonmetals are called the halogens – from the
Greek halos meaning salt and genesis meaning formation
The halogens are extremely reactive and react readily with
metals. Within the halogen group, fluorine is the most
reactive and reactivity decreases as we move down the
group
Group VIIIA are called the noble gases, which are extremely
UNreactive
The elements that border the staircase are the
metalloids (only element touching the staircase
that is not a metalloid is aluminum)
 Metalloids have properties that can be similar to
both metals and nonmetals, depending on
temperature, other compounds present, etc.
 Silicon is a good example. In a computer, silicon is
used to conduct electricity in circuit boards. In
baking, silicon is used to make heat resistant oven
mitts
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Classification of the Elements
SECTION 6.2
ORGANIZATION BY ELECTRON CONFIGURATION
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What is a valence electron?
Atoms in the same group have the same number of
valence electrons
Because elements in the same group have the same
number of valence electrons, they have similar chemical
properties
Within a period, the number of valence electrons
increases from left to right
The number of valence electrons in representative
elements can be found on the periodic table. The
number paired with the A heading indicates the number
of valence electrons
THE S-, P-, D-, AND F-BLOCK ELEMENTS

Remember electron configurations? Write the
electron configuration for iron
The electron dot diagram explains why we start
a new row after each noble gas
 Their outer electron level gets full, so anything
added after needs to be added to a new energy
level
 The rows on the periodic table indicate what
energy level those electrons are occupying

The periodic table is divided into blocks, the s-,
p-, d-, and f-block
 These blocks are the same as the s, p, d, and f
energy sublevels we talked about earlier
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Periodic Trends
SECTION 6.3
ATOMIC RADIUS
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Technically defined, atomic radius is the area where
there is a 90% probability of finding electrons
Essentially, it is the size of the atom
Atomic radius increases from top to bottom within a
group. Why?
Within a period, atomic radius decreases. Why?
It decreases because we are adding protons and
electrons, but staying on the same energy level. This
increases effective nuclear charge which pulls the
electrons in tighter, making the atom smaller
IONIC RADIUS
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Atoms can gain or lose electrons
When they do this they form an ion, which is an atom
that has a net positive or negative charge
When an atom loses electrons, it forms a positive charge
and the radius decreases from its neutral atom
When an atom gains electrons, it forms a negative
charge and the radius increases from its neutral atom
Within a period ionic radius of positive ions decreases,
then increases as it changes from positive to negative
ions, then decreases again
Within a group, ionic radius increases
IONIZATION ENERGY
Ionization energy – the energy required to
remove an electron from an atom
 Ionization energy increases from left to right
across a row. Why?
 Ionization energy decreases from top to
bottom. Why?
 Atoms can have second or third ionization
energies as well
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ELECTRONEGATIVITY
Electronegativity is the desire for electrons
 Electronegativity increases from left to right
across a period
 Electronegativity decreases from top to bottom
 The noble gases have very minimal
electronegativity
 The most electronegative element is fluorine,
while the least is francium
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