Transcript Trends in the Periodic Table

Trends in the Periodic Table
•A trend is a predictable change in a
particular direction.
•Example: There is a trend in the alkali
metals to increase in reactivity as you
move down a group.
Atomic Radius
• Bond radius: one half the distance
between the nuclei of two identical atoms
that are bonded together.
• The outer edge of the atom is not easy to
determine because the path of electrons is
not well-defined.
Group Trend
• The atomic radius increases as you move down
a group
• Adding electrons to an additional energy level
with each period
• Electron shielding - the reduction of the
attractive force between a positively charged
nucleus and its outermost electrons due to the
cancellation of some of the positive charge by
the negative charges of the inner electrons
Period Trend
• Atomic radius decreases as you move across
the period.
• Electrons are being added to the same energy
level, (there is no electron shielding) the
negative charge is increasing with each electron
added.
• The positive charge of the nucleus is increasing
with each element in the period.
• The nucleus is pulling the electrons closer.
• At the point where atomic radius levels off the
electrons begin repelling each other because
they are close together.
Ionization Energy
• The energy required to remove an electron from
an atom
• The process can be expressed as
A + ionization energy  A+ + e• Ionization energy increases going across a
period
• The attraction for outer shell electrons becomes
greater as the pos. nuclear charge increases
with a decrease in atomic radius.
• Atoms are becoming more stable with half filled
or approaching the completely filled state.
Ionization Energy Group Trends
• Ionization energy decreases going down a
group
• Atoms are larger and outer shell electrons
are farther from the nucleus.
• Electron shielding reduces the attraction
for outer shell electrons by the nucleus.
Additional Ionization Energies
• With each additional ion removed the
ionization energy will increase.
• You will see large increases in ionization
energies when a noble gas configuration is
reached.
• Example Li+ has the same electron
configuration as He so you will see a large
jump between the first and second
ionization energies.
Electronegativity
• The ability of an element in a compound to
attract electrons.
• Linus Pauling is the chemist who made the
scale.
• Fluorine is the most electronegative of all
elements and has been assigned the
value of 4.
• All other elements are assigned values in
relationship to fluorine.
Group trends
• Within a group the electronegativity
decreases as you move down a group
• The nucleus has less attraction for outer
shell electrons as the atom gets larger and
because of electron shielding
Period trends
• Electronegativity increases going across a
period
• Increasing nuclear charge is the main reason for
the increase
• The increase across the period is more
pronounced than the decrease down a group.
Electron shielding causes the effective nuclear
charge to about the same going down a group
so the main difference is the increase in distance
between the nucleus and the outer shell
electrons
Electron Affinity
• The energy change associated with the gain of
an electron by an individual atom.
• A positive value mean energy is required to gain
an electron. Elements with positive values do
not gain electrons easily and will most likely lose
the electron once the energy is removed.
• A negative value means energy is given off
when an electron is gained. The more energy
given off the easier it is to gain an electron.
Elements with high negative electron affinity
values gain electrons easily.
Group and Period Trends
• Electron Affinity decreases going down a
group.
• Electron Affinity increases going across a
period
• The reasons are the same as for
electronegativity even though these are
different properties.
Ionic Radius
• Atoms form ions by either losing or gaining
electrons
• Metals tend to lose electrons and
Nonmetals tend to gain electrons
• The ion size of metals is smaller than the
neutral atom
• The ion size of nonmetals is larger than
the neutral atom
Group and Period Trends
• Ions tend to get larger as you move down
a group just as neutral atoms do.
• Ions tend to get smaller moving across a
period until anions (negative ions) begin
forming rather than cations (positive ions).
• There is an increase where anions begin
then there is a general decrease
Melting and Boiling Points
• Increase going across a period until the
orbitals are half full then there is a
decrease until the orbitals are completely
full at which point the melting and boiling
points begin increasing again.