Note: Studying pp. 120 - 126 in the textbook is also strongly encouraged.
Note: These are general periodic trends of elements. There are
many exceptions to these general rules.
Review
Period - a row of elements on the
periodic table. Remember that sentences are written in rows and end with
a period.
Group - a column of elements on
the periodic table. Remember that group is spelled group and groups go
up and down.
Atomic Radius
- Atomic radius is simply the radius of the atom, an indication of the
atom's volume.
Period - atomic radius decreases
as you go from left to right across a period.
Why? Stronger attractive forces in atoms
(as you go from left to right) between the opposite charges in the nucleus
and electron cloud cause the atom to be 'sucked' together a little tighter.
Group - atomic radius increases as
you go down a group.
Why? There is a significant jump in the
size of the nucleus (protons + neutrons) each time you move from period
to period down a group. Additionally, new energy levels of elections clouds
are added to the atom as you move from period to period down a group, making
the each atom significantly more massive, both is mass and volume.
Electronegativity
- Electronegativity is an atom's 'desire' to grab another atom's electrons.
Period - electronegativity increases
as you go from left to right across a period.
Why? Elements on the left of the period
table have 1 -2 valence electrons and would rather give those few valence
electrons away (to achieve the octet in a lower energy level) than grab
another atom's electrons. As a result, they have low electronegativity.
Elements on the right side of the period table only need a few electrons
to complete the octet, so they have strong desire to grab another atom's
electrons.
Group - electronegativity decreases
as you go down a group.
Why? Elements near the top of the period
table have few electrons to begin with; every electron is a big deal. They
have a stronger desire to acquire more electrons. Elements near the bottom
of the chart have so many electrons that loosing or acquiring an electron
is not as big a deal. This is due to the shielding affect where electrons
in lower energy levels shield the positive charge of the nucleus from outer
electrons resulting in those outer electrons not being as tightly bound
to the atom.
Ionization Energy
- Ionization energy is the amount of energy required to remove the outmost
electron. It is closely related to electronegativity.
Period - ionization energy increases
as you go from left to right across a period.
Why? Elements on the right of the chart
want to take others atom's electron (not given them up) because they are
close to achieving the octet. The means it will require more energy to
remove the outer most electron. Elements on the left of the chart would
prefer to give up their electrons so it is easy to remove them, requiring
less energy (low ionization energy).
Group - ionization energy decreases
as you go down a group.
Why? The shielding affect makes it easier
to remove the outer most electrons from those atoms that have many electrons
(those near the bottom of the chart).
Reactivity - Reactivity
refers to how likely or vigorously an atom is to react with other substances.
This is usually determined by how easily electrons can be removed (ionization
energy) and how badly they want to take other atom's electrons (electronegativity)
because it is the transfer/interaction of electrons that is the basis of
chemical reactions.
Metals
Period - reactivity decreases as
you go from left to right across a period.
Group - reactivity increases as
you go down a group
Why? The farther to the left and down
the periodic chart you go, the easier it is for electrons to be given or
taken away, resulting in higher reactivity.
Non-metals
Period - reactivity increases as
you go from the left to the right across a period.
Group - reactivity decreases as
you go down the group.
Why? The farther right and up you go on
the periodic table, the higher the electronegativity, resulting in a more
vigorous exchange of electron.
Ionic Radius vs. Atomic Radius
Metals - the atomic radius of a metal is generally larger than the
ionic radius of the same element.
Why? Generally, metals loose electrons
to achieve the octet. This creates a larger positive charge in the nucleus
than the negative charge in the electron cloud, causing the electron cloud
to be drawn a little closer to the nucleus as an ion.
Non-metals - the atomic radius of a non-metal is generally smaller than
the ionic radius of the same element.
Why? Generally, non-metals loose electrons
to achieve the octet. This creates a larger negative charge in the electron
cloud than positive charge in the nucleus, causing the electron cloud to
'puff out' a little bit as an ion.
Melting Point
Metals - the melting point for metals generally decreases as you go
down a group.
Non-metals - the melting point for non-metals generally increases as
you go down a group.
