Chapter 6
The Periodic Table and Periodic Law
I. History of the Periodic Table
A. Just a list
• In the late 1790s, there were only 23
known elements
• The advent of electricity made it possible
to break down compounds into their
component elements
http://www.youtube.com/watch?v=OTEX38bQ-2w
• The development of the spectrometer made it
possible to identify new elements
• By 1870 there were 70 known elements
• Scientists needed a tool for organizing the
many facts associated with the elements
B. John Newlands
• Noticed that when the elements were
arranged by increasing atomic mass, their
properties repeated every eighth element
• A pattern such as this is called periodic
C. Mendeleev and
Moseley
• Mendeleev organized
the elements into the
first periodic table
• He predicted the
existence and
properties of
undiscovered
elements
• Moseley rearranged the
elements by increasing
mass number, or number
of protons, for a more
accurate table
• The repetition of chemical
and physical properties of
elements by increasing
atomic number is called
the periodic law
II. The Modern Periodic Table
• Elements are arranged in order of
increasing atomic number into a series of
columns, called groups or families, and
rows, called periods
• Each group is number 1 through 8, followed
by the letter A or B
- the groups designated with an A are often
referred to as the main group, or
representative elements
- the groups designated with a B are referred
to as the transition elements
A. Classifying the elements
1. Metals
• Generally shiny when smooth and clean
• Solid at room temperature
• Good conductors of heat and electricity
• Ductile, or easily pounded into thin
sheets
• Malleable, or easily drawn into wires
a. Alkali metals
• Group 1A elements (except for Hydrogen)
• Extremely reactive
b. Alkali earth metals
• Group 2A elements
• Less reactive than the alkali metals
c. Transition metals
d. Inner transition metals
i. Lanthanide series
• Used extensively as phosphors,
substances that emit light when struck
by electrons
ii. Actinide series
2. Nonmetals
• Generally gases or brittle, dull-looking
solids
• Poor conductors of heat and electricity
• The only liquid nonmetal is bromine (Br)
a. Halogens
• Group 7A
• Highly reactive
b. Noble gases
• Group 8A
• Extremely unreactive
3. Metalloids, or semimetals
• Elements with physical and chemical
properties of both metals and nonmetals
III. Organizing the Elements by Electron
Configuration
A. Valence electrons
• Electrons in the highest principal energy
level
• Atoms in the same group have similar
chemical properties because they have
the same number of valence electrons
B. Valence electrons and period
• The energy level of an element’s valence
electrons indicates the period on the
periodic table in which it is found
ex. Lithium’s valence electron is in the
second energy level and lithium is found
in period 2
C. Valence electrons and group number
• A representative element’s group number
indicates the number of valence electrons it
has (with the exception of helium)
IV. The s-, p-, d-, and f-block Elements
• The periodic table is divided into four
blocks
A. s-block elements
• Consists of groups 1A and
2A as well as helium
• Valence electrons occupy
only s orbitals
• Group 1A elements have
partially filled s orbitals
containing one valence
electron (s1)
• Group 2A elements have
completely filled s orbitals
containing two valence
electrons (s2)
• Because s orbitals hold a
maximum of two electrons,
the s-block portion of the
periodic table spans two
groups
B. p-block elements
• After the s-block is
full, valence electrons
next occupy the p
sublevel and its three
p orbitals
• The p-block spans six
groups on the periodic
table because three p
orbitals can hold a
maximum of six
electrons
• The group 8A elements have
s and p sublevels that are
completely filled
- this results in an unusually
stable atomic structure
leaving the elements
virtually unreactive
C. d-block elements
• Contains the transition metals
• The five d orbitals can hold a total of ten
elements; thus the d-block spans ten
groups on the periodic table
• Have a filled outermost s orbital of energy
level n
• Have filled or partially filled d orbitals of
energy level n-1
ex. Titanium = [Ar] 4s23d2
s-orbital =
n
=4
d-orbital =
n-1 = 3
D. f-block elements
• Contains the inner transition metals
• Because there are seven f orbitals holding
up to a maximum of 14 electrons, the fblock spans 14 columns
E. Putting it all together
• As you proceed down through the periods, the
principle energy level increases as well as the
number of energy sublevels
Period 1: s-block elements
Periods 3 and 4: s- and p-block elements
Periods 4 and 5: s-, p-, and d-block elements
Periods 6 and 7: s-, p-, d-, and f-block elements
V. Atomic Radius
• The outer limit
of an electron
cloud is defined
as the spherical
surface within
which there is a
90% probability
of finding an
electron
• Atomic size is defined by how closely an
atom lies to a neighboring atom, which
varies from element to element
A. Trends within periods
• Atomic size decreases left-to-right across a
period
• Each successive element has one additional
electron in the same principal energy level
• Each element also gains one more proton
that pulls the outermost electrons closer to
the nucleus
B. Trends within groups
• Atomic size increase as you move down a
group
• A principal energy level is added to
elements in each period going down the
table
• Each orbital also increases in size,
separating valence electrons further from
the nucleus which offsets the pull of the
increased nuclear charge
Atomic Radius
VI. Ionic Radius
• An ion is an atom that has a positive or
negative charge
A. When atoms lose electrons and form
positively charged ions, they always
become smaller
• The electron lost is always a valence electron
1. may leave an empty outer orbital
2. the repulsion between electrons is less,
allowing them to be pulled closer to the
nucleus
B. When atoms gain electrons and form
negatively charged ions, they always become
larger
• The electrostatic repulsion between the
atom’s outer electrons forces them to move
apart
C. Trends within periods
1. Positive ions on the left side of the table
become smaller left-to-right
2. Negative ions on the right side of the table
become smaller left to right
D. Trends within groups
• As you move down a group, an ion’s outer
electrons are in higher principal energy
levels, resulting in a gradual increase in
ionic size
VII. Ionization Energy
• Ionization energy is the energy required to
remove an electron from an atom
• Energy is needed to overcome the
attraction between the positive charge in
the nucleus and the negative charge in the
electrons
- a high ionization energy indicated the
atom has a strong hold on its electrons
- a low ionization energy indicates an
atom loses its outer electrons easily
• Group 1A elements have low ionization
energies
• Group 8A elements have high ionization
energies
• The energy required to remove the first
electron from an atom is called the first
ionization energy
• The energy required to remove the second
electron is called the second ionization energy
• For each element there is an ionization for
which the required energy jumps dramatically
- related to the atom’s number of valence
electrons
- atom’s hold onto their core electrons much
more strongly than they hold onto their
valence electrons
A. Trends within periods
• Generally increases left-to-right
B. Trends within groups
• Generally decreases down a group
• Valence electrons farther from the nucleus
require less energy to remove them
C. Octet rule
• States that atoms tend to gain, lose, or
share electrons in order to acquire a full set
of eight valence electrons
- hydrogen and helium are complete with
two valence electrons
1. Elements on the right side of the table tend
to gain electrons (become negative ions) to
acquire the noble gas configuration
2. Elements on the left side of the table tend to
lose electrons (become positive ions)
VIII. Electronegativity
• Indicates the relative ability of its atoms
to attract electrons in a chemical bond
• Expressed in Paulings; numerical values
of 4.0 or less
• Fluorine is the most electronegative
element (3.98)
• Cesium (0.79) and Francium (0.7) are the
least electronegative elements
A. Trends within periods and groups
• Generally decreases as you move down a
group
• Increases from left-to-right