To understand the trends of the periodic table, they must understand effective nuclear charge.
Effective nuclear charge is the attractive force that a nucleus has for electrons.
ATOMIC RADIUS – the size of an atom (the distance between the nucleus and valence electrons).
The atomic radius of an atom is usually measured in picometres.
Down a group: Atomic Radius
As you go down a group...
as you move down the group.
Valence electron experiences...
Across a period: Atomic Radius
As you go down a group...
from left to right.
Valence electron experiences...
Figure 1: Periodic trend of atomic radii. Atomic radii
increases as you move towards the left and downwards
in the periodic table.
IONIC RADIUS – the size of an ion (the distance between the nucleus of the ion and the valence
electrons)
FOR METAL CATIONS
The ionic radius is
for a metal ion than the atomic radius of the
metal atom
Consider what would happen if an alkali
metal lost one electon...

electrons (p+>e-)
 There are
shells
 Valence electrons are
to the nucleus
 The attractive force of the nucleus is
greater and is shared between fewer
electrons (ENC
)
FOR NON-METAL ANIONS
The ionic radius is
for a non-metal ion than the atomic radius of
the non-metal atom.
Consider what would happen if a non-metal
gained one electron...

(e->p+)
 There are the
number of
shells
 More valence electrons means electron

The charge of the nucleus is the same
and the ENC is shared between more
electrons (ENC
)
Figure 2: Periodic trend of ionic radii for metal cations and non-metal anions
IONIZATION ENERGY – the energy required to remove an electron from its outer shell.
Ionization energy is measured in kilojoules per mole (kJ/mol).
Down a group: Ionization Energy
REASONS:
as you move down the group.
Across a period: Ionization Energy
REASONS:
from left to right across a period.
What about noble gases?
Discuss the trend in
Ionization energy with
your table partners and
justify your reasoning
using ENC.
Figure 3: A chart representing the periodic trend of ionization energy. Ionization energy increases as
you move towards the right and downwards in the periodic table.
ELECTRON AFFINITY – the energy release that occurs when an electron is added to a neutral
atom, Electron affinity is measured in kilojoules per mole (kJ/mol).
ELECTRONEGATIVITY – attraction of an atom for an electron within a chemical bond.
Down a group: Electron Affinity and Electronegativity
down the group.
REASONS:
as you move
Across a period: Electron Affinity and Electronegativity
right across a period, especially for the Noble Gases.
REASONS:
from left to
Figure 4: Various diagrams showing how the size of the atom affects its abilities to attract electrons.
REACTIVITY – the ability of an atom to react
Starting from the middle, reactivity
moving outwards to the ends of the periodic table (except for
Noble Gases)
REASON: Moving outward, atoms need to give fewer electrons to
achieve an octet
On the metal side, reactivity
moving down the
group and right to left
REASON: Large metals give up electrons easier, as they are
farther away from nucleus
On the non-metal side, reactivity
moving up
the group.
REASON: Smaller non-metals hold the incoming electrons more
strongly
Figure 5: Reactivity of Metals
Figure 6: Summary of Periodic Trends including atomic radius, electron affinity, ionization energy and
reactivity of metals (metallic character) and reactivity of non-metals (non-metallic character)
HOMEWORK: Read Pages 36-41 and do questions #1-9 on Page 41. Try Question 10.