Periodic Table & Periodicity
Ms Piela
Durfee High
Periodic Trends/ Periodicity
 A periodic trend is a pattern observed
on the periodic table for an atomic
property
 Each of the four trends have
explanations for their group trend and
their period trend
The 4 Main Periodic Trends are:
Atomic Radius
Ionization Energy
Electronegativity
Electron Affinity
The Period Trend Explanation
 When comparing elements in the
same period, compare the effective
nuclear charges (Zeff)
 Effective nuclear charge is the net positive
charge experienced by electrons in an atom
The Period Trend Explanation
 The atoms on the right of the periodic
table have higher effective nuclear charges
(Zeff) when compared to elements on the
left
 This is due to electrons being added to the same
energy level. They are approximately the same
distance away from the nucleus
 In general, the further atoms are away from the
nucleus, the less attracted they become
The Group Trend Explanation
 When comparing atoms in the same
group, compare the amount of
electron shielding occuring
 Electron shielding is where core electrons
shield outer electrons from the charge of the
nucleus
 Thus, outer electrons are held less tightly
because of electron/electron repulsion
The Group Trend Explanation
 Atoms on the top of the periodic table have less
electron shielding than atoms at the bottom
 As you increase in the number of energy levels, more electron shielding
occurs
 This does NOT occur across a period as energy levels will not change
Atomic Radius
 Atomic Radius is a measure of the size of the
atom
 Measured by the distance from the nucleus to the outermost
electrons
Atomic Radius
 Atomic Radii decreases moving across a
period, and increases going down a group
 For the period trend: with effective nuclear charge,
the increased positive charge pulls electrons closer,
causing the size to decrease
 With the group trend, the increasing energy levels
provide more electrons, which increase the size of
the atom (electron shielding doesn’t really work)
Ionization Energy
 Ionization energy is the energy
required to remove an electron from
an atom
 Amount of energy increases as the number of
ionizations occur (i.e. first ionization takes less
energy than the second, and so on)
Ionization Energy
 Ionization energy increases going across a
period and decrease going down a group
 With increasing effective nuclear charge, electrons
are held more tightly, thus atoms on the right
require more energy to remove an electron
 With increasing electron shielding, electrons are held
less tightly and thus decrease in IE
Graph of IE Periodic Trend
Electronegativity
 Electronegativity is the ability of an atom in
a molecule to attract shared electrons to
itself
 Think of electronegativity as a “tug of war”
Electronegativity
 Electronegativity increases going across a
period, and decreases going down a group
 Due to increasing effective nuclear charge, atoms on
the right hold electrons more tightly, causing them
to have high EN
 Due to electron shielding, atoms on thebottom tend
to hold electrons more loosely, making them have
low EN
Electronegativity
 The noble gases are excluded from this trend as they
tend not to bond with other atoms
 This makes fluorine the most electronegative atom
Electron Affinity
 Electron Affinity is the energy
associated with the addition of an
electron to an atom
 The more negative the quantity, the more
energy is released upon the addition of an
electron
Electron Affinity
 Electron affinity increases across a period
and decreases going down a group
 Due to increasing effective nuclear charge, atoms on
the right tend to want to attract negative electrons
more
 Due to electron shielding, atoms on the bottom tend
to hold electrons more loosely, making them have
low EA