Leaving Certificate Chemistry
Chapter 7 – Trends in the Periodic Table
Atomic radii
The atomic radius of an atom is defined as half
the distance between the nuclei of two atoms
of the same element that are joined together
by a single covalent bond.
H
Li
.123
Be
.0
89
B
..08
0
C
.07
7
N
.07
4
O
.0
74
F
.072
Na
.157
K
.203
Rb
0.8
Cs
0.7
Atomic radii decrease across the table
because:
1. Increasing nuclear charge (more
protons) pulls the electrons closer to
the nucleus, and
2. The screening effect of inner electron
shells will be the same for all members
of a given period. The combined effect
of both factors results is the electrons
being pulled closer to the nucleus and a
smaller radius.
Atomic radii increase down a group because,
1. A new shell is added for each
successive member of a group, leading
to a greater radius.
2. Increased screening effect of extra
electron shells i.e. the nucleus has less
of a pull on the outer electrons.
Ionisation energies
The 1st ionisation energy of an element is
the energy needed to remove the most
loosely bound electron from a neutral
gaseous atom of that element.
The process may be represented by the
following equation,
X  X   e
Ionisation energy values increase across the
table because:
1. Increasing nuclear charge - bigger pull
on the electrons, harder to remove.
2. Decreasing atomic radius - electrons
nearer to nucleus, bigger pull on them.
Harder to remove.
ionisation energies of
a period
2500
2000
1500
1000
500
0
Li Li Be B C N O F Ne
A quick glance, at the graph of 1st ionization
energies of period 2, reveals unusually high
values for Be and N. The reasons for this
become evident if we look at the electronic
configurations of these two elements:
Be; 1s2, 2s2
Beryllium has a full 2s orbital which makes it
quite stable, therefore removing an electron
from this element takes more energy than
expected.
Leaving Certificate Chemistry
N; 1s2,2s2,2px1,2py1,2pz1
Nitrogen has three half-filled 2p orbitals, this
makes it very stable and removal of an electron
is more difficult than expected.
Ionisation energy values decrease down a
group because:
1. Increasing atomic radius - electrons
further from nucleus, nuclear pull is
less, easier to remove.
2. An extra shell is added for each
successive member of a group,
therefore the screening effect is
greater as you go down a group. The
outer electrons are further removed
and more easily removed.
Subsequent ionisations
If you study the 1st and subsequent ionisation
energy values of an element such as potassium,
you will observe some significant jumps in the
energy values as you remove more and more
electrons. You will see a general increase as you
remove each subsequent electron.
120000
100000
ionisation
energy
80000
60000
40000
20000
11
th
13
th
15
th
17
th
19
th
9t
h
7t
h
5t
h
3r
d
1s
t
0
electrons removed
But, more significantly, huge jumps occur
when removing the 2nd , 10th , and 18th
electrons, as to do so means breaking into a
completely full energy level. This is further
evidence for the existence of energy levels.
Trends in electronegativity values
‘The electronegativity of an element is a
measure of the power of attraction, of an
atom of that element, for the shared pair
of electrons in a covalent bond.’
Electronegativity values increase across the
table because:
1. The nuclear charge is increasing
and
2. The atomic radius is decreasing.
These two factors combined make it more
difficult to remove an electron.
Electronegativity values decrease down a
group because:
1. Increasing atomic radius
2. Increased screening effect.