Transcript Periodic Table

Periodic Table
Periodic Table
 Placed
in order of their atomic numbers
 The similar elements are placed in
columns, known as groups or families
 The rows of the table are called periods
How did Mendeleev’s periodic table
differ from the modern periodic
table?
The elements in Mendeleev’s table were listed in
order of atomic weights rather than atomic
numbers. Atomic numbers were unknown in
1871. None of the internal structure of the atom
was known in 1871.
 Noble gases are completely absent--they were
not known in 1871. The first noble gas
discovered in 1898 was argon.

Representative Elements

Main Group Elements – Outer subshell is an s
and p in the highest energy level
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
Group 1a – Alkali
Group 2a – Alkaline Earth
Group 3a – Boron Family
Group 4a – Carbon Family
Group 5a – Nitrogen Family
Group 6a – Oxygen Family
Group 7a – Halogens
Group 8a – Noble Gases
Write the final energy level for each group
Groups 1b to 8b
 Transition

Metals
Incompletely filled d subshell
Valence Electrons
 Outer
electrons of an atom, which are
involved in bonding
Ionization Energy
 Ionization
energy is defined as the energy
required to remove an electron from a
gaseous atom.
X(g) + energy → X+(g) + e-
 For
every row of the periodic table, the
ionization energy increases from left to
right (with some exceptions) reaching a
maximum on the noble gases. It then
starts low again for the next alkali metal
(Group IA) and gradually increases again
reaching a maximum on the next noble
gas.
The periodic trends are:
Horizontal: IE increases across a row of the
periodic table (from left to right)
Vertical: IE decreases down a group
(column) of the periodic table (from top to
bottom)
What are the reasons for these
trends?


Horizontal: As you go across from left to right in a row of
the periodic table, you are adding electrons to the same
shell but with increasing nuclear charge. The increasing
number of protons (higher Z) attracts the electrons more,
making it harder to remove an electron from the atom-hence a higher IE.
Vertical: As you go down a group from top to bottom,
you always have the same valence shell configuration.
However, each succeeding shell is further from the
nucleus, and is shielded from the nuclear pull by inner
electrons. It is thus easier to remove electrons from
outer shells which are less attracted to the nucleus.
Atomic Radius
 Atomic
radii measure the size of the atom.
Although the atom does not have a distinct
boundary, there are several ways to
estimate atomic radii based on distances
between atoms in crystals or molecules.
Atomic radii show distinct trends.
 Horizontal:
atomic radius decreases
across a row of the periodic table (from left
to right)
 Vertical:
atomic radius increases down a
group (column) of the periodic table (from
top to bottom)
What are the reasons for these
trends?
 Horizontal:
As you go across from left to
right in a row of the periodic table, you are
adding electrons to the same shell but with
increasing nuclear charge. The increasing
number of protons (higher Z) attracts the
electrons more, making it the electron
cloud closer to the nucleus; hence a
smaller atomic radius.
 Vertical:
As you go down a group from top
to bottom, you always have the same
valence shell configuration. However,
each succeeding shell is further from the
nucleus, and is shielded from the nuclear
pull by inner electrons. So the outer shell
logically is larger there are more shells
(higher principal quantum number).
Most atoms form ions. Which is
larger, the ion or the atom from
which it is formed?
 Positive
ions are formed when the atom
gives off an electron. The resultant ion is
always smaller than the corresponding
atom, since the resultant positive charge
(more protons than electrons) causes the
remaining electron cloud to be pulled
toward the nucleus.
 Negative
ions are formed when the atom
gains an electron. The resultant ion is
always larger than the corresponding
atom, since the resultant negative charge
(more electrons than protons) causes the
electron cloud to be repelled away from
the nucleus.
Electron Affinity
 When
an electron is added to a gaseous
atom, forming a negative ion, energy may
be either released or absorbed.
 The ΔH for this process is defined as the
electron affinity.
X(g) + e- → X- (g)
ΔH = electron affinity of X
What are the trends in electron
affinity?
 Horizontal:
Going from left to right across
a row, the electron affinity gets more
negative (more attraction for electrons),
with the halogens (not the noble gases)
having the most negative electron affinity
(most attraction for electrons).
 Vertical: Going from top to bottom down a
group, the electron affinity gets less
negative (less attraction for electrons.
What are the reasons for these
trends?


Horizontal: As you go across from left
to right in a row of the periodic table, you
are adding electrons to the same shell but
with increasing nuclear charge. The
increasing number of protons (higher Z)
attracts the electrons more; a higher
attraction for electrons means a more
negative electron affinity.
 Vertical:
As you go down a group from top
to bottom, you always have the same
valence shell configuration. However,
each succeeding shell is further from the
nucleus, and is shielded from the nuclear
pull by inner electrons. Less attraction by
the nucleus means a less negative
electron affinity.