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Chapter 4
The Periodic table
Mendeleev’s Table
 Mendeleev’s table grouped elements with
similar properties into vertical columns
called “groups” or “families” example the
halogen group.
 Mendeleev found only 63 elements known at
that time. There are 2 interesting things to note.
 The elements do not always fit neatly in the
order of increasing atomic mass.
 There were gaps in his table where elements
with a particular atomic mass should occur. He
was able to predict the properties of these
undiscovered elements which he called
ekaaluminium , ekaboron and ekasilicon.
 Eka meaning “one beyond” in Sanskrit.
Moseley’s theory
 Mosley an England scientist worked with x-ray spectra
to study atomic structure which helped him determine
atomic numbers for chemical elements. He discovered
isotopes, explaining how atomic mass did not order the
elements appropriately. Moseley was the catalyst for the
periodic law, a rule stating that the properties of
elements are a periodic function of their atomic numbers.
Using the atomic numbers, He was able to more
accurately position the elements in the periodic table.
Moseley's change to the periodic table allowed the few
problems with Mendeleev's periodic table to disappear.
The modern periodic table is now based on atomic
 We now follow the periodic table
recommended by IUPAC.
 The main group elements:
 Groups 1,2, and 13 through 18 are
referred to as the main group elements.
 Group 1 the alkali metals. These metals
react with water to form alkaline solution.
They are excellent conductors of
 Group 2- The alkaline earth metals. These
are harder than alkali metals. They are
less reactive than alkali metals. Mg from
this group when reacts with oxygen from
the air forms magnesium oxide which
slows corrosion of the magnesium metals
underneath. Forms good alloys.
 Lanthanides and actinides: the lanthanides have atomic
numbers from 58 to 71.The name originated from the
element 57. they are shiny, reactive metals that have
irregular electronic configuration.
 Actinides are unique in that their nuclear structures are
of more importance than their electron configurations.
They have an unstable arrangement of protons and
neutrons in the nucleus so they are called radioactive.
Example Uranium. Nuclear disintegration of the uranium
nucleus releases sufficient energy to run power
parts,submarines also aircrafts.
 Group 17-Halogens Derived from the
Greek word meaning “salt former”. Most
reactive group of nonmetals.
 Group 18- Noble gases. Formerly called
inert gases because they were thought to
be completely unreactive. Except for
helium these atoms are characterized by
an octet of electrons, ns2 and np6 in the
outermost energy level.
 Page 123 # 7,9,11,13
Characterestics Of metals
 Metals are good conductors of heat and
electricity. Metals in the form of crystals .
 A crystal is a substance in which atoms or
molecules are arranged in an orderly
geometric fashion.
Crystal and conduction band
 The outer electrons in the atoms of each
element form the bonds that bind the
atoms together in a crystal. In a crystal,
each atom is bonded to all neighboring
atoms. These electrons exist in orbitals
just as electrons do in atoms. There are
vacant orbitals in the bond just as in
atoms. Electrons are free to move in three
dimensions through the crystal, forming
the conduction band.
Conduction band is a band within which
electrons must move to allow electrical
conduction. In metals the orbitals of the
conduction band actually overlap
and are continous with
the bond orbitals so even
at room temp. there
are many electrons in
the conduction band.
 The metal atoms are vibrating back and
forth with thermal motion, interfering with
the motion of electrons. The vibration
increases as the temperature increases so
the conductivity decreases.
 In semiconductors the energy of the higher
orbitals is not very high compared with the
energy of the filled bonding orbitals.less
energy is required to excite electrons to
these orbitals, so even at room
temperature they carry a certain amount of
current. At higher temperature there are
more electrons excited in the conducting
band, so conductivity of semiconductors
increases as the temperature increases.
 In typical non metals these orbitals have
very high energy so it is practically
impossible to excite electrons to the
orbitals of the conduction band example
diamond is a nonconductor.
Trends in periodic table
 Trend is a periodic change in a particular
 Periodic trends in atomic radii: The size of
an atom is determined by the volume
occupied by the electrons surrounding the
nucleus. Radii are usually determined for
atoms that are chemically bonded or close
together in the solid state, as in Bromine
and Iodine.
 Bond radii: Half the distance between the
nuclei of the atoms in each molecule is the
bond radius.
 The distance between the
nuclei in adjacent non
bonded molecules, which
is twice the distance, called
the van der Waals radius.
 Van der waal’s radii are not very precise
because of the fuzziness of the atoms.
Bond radius is usually considered to dfind
the size of the atom.
 Atomic radius increases as you move
down a group. This is caused by the
addition of another principal energy level
from one period to the next.
 Electron shielding: The electrons in the
inner energy levels are between the
nucleus and the outermost electrons and
therefore shield the outer electrons from
the full charge of the nucleus
 Atomic radius decreases as you move across a period.
 As you move from left to right across a period, each
atom has one more proton and one more electron than
the atom before it has.
All additional electrons go into the same principal energy
level—no electrons are being added to the inner levels.
 Electron shielding does not play a role as you move
across a period.
 As the nuclear charge increases across a period, the
effective nuclear charge acting on the outer electrons
also increases.
 Ionization Energy: The energy used to remove the
electron is the ionization energy of the atom.
 Each element has more occupied energy levels than the
one above it has.
 The outermost electrons are farthest from the nucleus
in elements near the bottom of a group.
 As you move down a group, each successive element
contains more electrons in the energy levels between the
nucleus and the outermost electrons.
 Ionization energy tends increases as you move from left
to right across a period.
 From one element to the next in a period, the number
of protons and the number of electrons increase by one
 The additional proton increases the nuclear charge.
 A higher nuclear charge more strongly attracts the outer
electrons in the same energy level, but the electronshielding effect from inner-level electrons remains the
 Electron affinity: The ability of an atom to
attract and hold an electron is called
electron affinity. Although the atom is
neutral the electrons in the orbitals
generally do not shield the nuclear charge
a 100%. The approaching electron may
experience a net pull because the effective
nuclear charge is greater than zero. The
electron enters a vacant orbital.
 For example, when a neutral chlorine atom
in the gaseous form picks up an electron
to form a Cl- ion, it releases an energy of
349 kJ/mol or 3.6 eV/atom. It is said to
have an electron affinity of -349 kJ/mol
and indicates that it
forms a stable negative ion
 Electron affinity becomes more negative across
a period and tends to decrease from top to
 Electron affinity trends within groups of elements
are not as regular as trends for ionization
energy. Alkaline earth metals are an exception.
An electron added in the alkaline earth metals
must go to the p orbital and is shielded to a
greater extent by the s electrons.
Electron affinity of the main block
elements in kj/mol
Periodic trends in melting and
boiling points
 As the number of electrons increases
across the period the m.p and b.p.
increases. This indicates stronger
However when the d orbitals become half filled
the m.p and b.p decreases. This indicates that
added electrons become less involved in
bonding. As the p electrons are added we see a
similar rise and fall with the maximum near the
stage where p orbitals are half filled. A minimum
is reached when the p orbitals are filled and the
atoms no longer bond. The noble gases are
monatomic and have no chemical bonding
forces between atoms. Hence their m.p. and b.p.
are unusually low.
 Page 151 Term review all
 Page 155 test prep all.