1.1.

10 Ionisation Energy
Ionisation Energies
The Ionisation Energy (IE) of an element is the amount of energy required to
remove one mole of electrons from one mole of gaseous atoms of an element to
form one mole of gaseous ions
Ionisation energies are measured under standard conditions which are 298 K and
101 kPa
The units of IE are kilojoules per mole (kJ mol-1)
The first ionisation energy (IE1) is the energy required to remove one mole of
electrons from one mole of atoms of an element to form one mole of 1+ ions
E.g. the first ionisation energy of gaseous calcium:
Ca(g) → Ca+ (g) + e- IE1 = +590 kJ mol-1
Trends in Ionisation Energies

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Ionisation energies show periodicity - a trend across a period of the Periodic

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Table
As could be expected from their electron configuration, the group 1 metals have a
relatively low ionisation energy, whereas the noble gases have very high ionisation
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energies
The size of the first ionisation energy is affected by four factors:
Size of the nuclear charge
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Distance of outer electrons from the nucleus
Shielding effect of inner electrons
Spin-pair repulsion

First ionisation energy increases across a period and decreases down a group
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A graph showing the ionisation energies of the elements hydrogen to sodium

Ionisation energy across a period
The ionisation energy across a period generally increases due to the following
factors:
Across a period the nuclear charge increases
This causes the atomic radius of the atoms to decrease, as the outer shell is
pulled closer to the nucleus, so the distance between the nucleus and the
outer electrons decreases
The shielding by inner shell electrons remain reasonably constant as electrons
are being added to the same shell
It becomes harder to remove an electron as you move across a period; more
energy is needed
So, the ionisation energy increases
Dips in the trend
There is a slight decrease in IE1 between beryllium and boron as the fifth electron
in boron is in the 2p subshell, which is further away from the nucleus than the 2s

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subshell of beryllium
Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron

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configuration is 1s2 2s2
Boron has a first ionisation energy of 800 kJ mol-1 as its electron
configuration is 1s2 2s2 2px1
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There is a slight decrease in IE1 between nitrogen and oxygen due to spin-pair
repulsion in the 2px orbital of oxygen
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Nitrogen has a first ionisation energy of 1400 kJ mol-1 as its electron
configuration is 1s2 2s2 2px1 2py1 2pz1
Oxygen has a first ionisation energy of 1310 kJ mol-1 as its electron
configuration is 1s2 2s2 2px2 2py1 2pz1
In oxygen, there are 2 electrons in the 2px orbital, so the repulsion between
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those electrons makes it slightly easier for one of those electrons to be

removed
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From one period to the next

There is a large decrease in ionisation energy between the last element in one
period, and the first element in the next period
This is because:
There is increased distance between the nucleus and the outer electrons as
you have added a new shell
There is increased shielding by inner electrons because of the added shell
These two factors outweigh the increased nuclear charge
Ionisation energy down a group
The ionisation energy down a group decreases due to the following factors:
The number of protons in the atom is increased, so the nuclear charge
increases
But, the atomic radius of the atoms increases as you are adding more shells of
electrons, making the atoms bigger
So, the distance between the nucleus and outer electron increases as you
descend the group
The shielding by inner shell electrons increases as there are more shells of
electrons
These factors outweigh the increased nuclear charge, meaning it becomes
easier to remove the outer electron as you descend a group
So, the ionisation energy decreases
Ionisation Energy Trends across a Period & going down a Group Table

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1.1.11 Ionisation Energy Trends
Ionisation Energies: Trends
Ionisation energies show periodicity - a trend across a period of the Periodic
Table
As could be expected from their electronic configuration, the group I metals have a
relatively low ionisation energy, whereas the noble gases have very high ionisation
energies
The size of the first ionisation energy is affected by four factors:
Size of the nuclear charge
The nuclear charge increases with increasing atomic number, which
means that there are greater attractive forces between the nucleus and
electrons, so more energy is required to overcome these attractive forces
when removing an electron

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Distance of outer electrons from the nucleus
Electrons in shells that are further away from the nucleus are less

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attracted to the nucleus - the nuclear attraction is weaker - so the further
the outer electron shell is from the nucleus, the lower the ionisation
energy
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Shielding effect of inner electrons
The shielding effect is when the electrons in full inner shells repel
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electrons in outer shells, preventing them from feeling the full nuclear
charge, so the more shells an atom has, the greater the shielding effect,
and the lower the ionisation energy
Spin-pair repulsion
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Electrons in the same atomic orbital in a subshell repel each other more
than electrons in different atomic orbitals which makes it easier to remove
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an electron (which is why the first ionization energy is always the lowest)
So, the first ionisation energy increases across a period and decreases down a
group
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A graph showing the ionisation energies of the elements hydrogen to sodium

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Ionisation energy across a period
The ionisation energy over a period increases due to the following factors:
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Across a period the nuclear charge increases
This causes the atomic radius of the atoms to decrease, as the outer shell is
pulled closer to the nucleus, so the distance between the nucleus and the
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outer electrons decreases
The shielding by inner shell electrons remain reasonably constant as electrons
are being added to the same shell
It becomes harder to remove an electron as you move across a period; more
energy is needed
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So, the ionisation energy increases

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There is a rapid decrease in ionisation energy between the last element in one
period, and the first element in the next period because:
There is increased distance between the nucleus and the outer electrons as
you have added a new shell
There is increased shielding by inner electrons because of the added shell
These two factors outweigh the increased nuclear charge
There is a slight decrease in IE1 between beryllium and boron as the fifth electron
in boron is in the 2p subshell, which is further away from the nucleus than the 2s
subshell of beryllium
Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron
configuration is 1s2 2s2
Boron has a first ionisation energy of 800 kJ mol-1 as its electron
configuration is 1s2 2s2 2px1
There is a slight decrease in IE1 between nitrogen and oxygen and phosphorus
due to spin-pair repulsion in the 2px orbital of oxygen
 Nitrogen has a first ionisation energy of 1400 kJ mol-1 as its electron
configuration is 1s2 2s2 2px1 2py1 2pz1
Oxygen has a first ionisation energy of 1310 kJ mol-1 as its electron
configuration is 1s2 2s2 2px2 2py1 2pz1
Ionisation energy down a group
The ionisation energy down a group decreases due to the following factors:
The number of protons in the atom is increased, so the nuclear charge
increases
But, the atomic radius of the atoms increases as you are adding more shells of
electrons, making the atoms bigger
So, the distance between the nucleus and outer electron increases as you
descend the group
The shielding by inner shell electrons increases as there are more shells of
electrons
These factors outweigh the increased nuclear charge, meaning it becomes

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easier to remove the outer electron as you descend a group
So, the ionisation energy decreases

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Ionisation energy trends across a period & going down a group table
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Successive ionisation energies of an element

The successive ionisation energies of an element increase
This is because once you have removed the outer electron from an atom, you have
formed a positive ion
Removing an electron from a positive ion is more difficult than from a neutral
atom
As more electrons are removed, the attractive forces increase due to decreasing
shielding and an increase in the proton to electron ratio
The increase in ionisation energy, however, is not constant and is dependent on
the atom's electronic configuration
Taking calcium as an example:
Ionisation energies of calcium table

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The first electron removed has a low IE1 as it is easily removed from the atom due
to the spin-pair repulsion of the electrons in the 4s orbital
The second electron is more difficult to remove than the first electron as there is
no spin-pair repulsion
 The third electron is much more difficult to remove than the second one
corresponding to the fact that the third electron is in a principal quantum shell
which is closer to the nucleus (3p)
Removal of the fourth electron is more difficult as the orbital is no longer full, and
there is less spin-pair repulsion

 Exam Tip
It is easy to remove electrons from a full subshell as they undergo spin-pair
repulsion.It gets more difficult to remove electrons from principal quantum
shells that get closer to the nucleus as there is less shielding and an
increase in attractive forces between the electrons and nuclear charge.

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