Ionisation Energy

Ionisation is the complete removal of an electron from an atom

or ion in the gaseous state.

Ionisation is an endothermic process as energy must be supplied

to the electron in order to overcome the electrostatic force

between the electron and the nucleus.

The energy required is called the standard ionisation energy.

The first ionisation energy is the heat enthalpy required to

remove the most loosely held electron from every atom in one

mole of free gaseous atoms under standard conditions.

The second ionisation energy is the energy required to remove

the most loosely held electron from every unipositive ions in

one mole of free gaseous unipositive ions under standard

conditions.

IONISATION ENERGY

Higher ionisation energies are defined in the same way. For

example, the third ionisation energy refers to:

However, the following is not the third ionisation energy:

The energy ∆HØ is the sum of the first, second and third

ionisation energy.

For any one element, the ionisation energies increase. As

each electron is removed from an atom, the remaining ion

becomes more positively charged. Moving the next electrons

away from the increased positive charge is more difficult and

the next ionisation energy is even larger.

IONISATION ENERGY

Factors Influencing Ionisation Energies

Generally, the bigger the size of an atom, the lower is the

ionisation energy.

The three strongest influences on ionisation energies of

elements are the following:

1) The Size of the Positive Nuclear Charge

This charge affects all the electrons in an atom. The

increase in nuclear charge with atomic number will tend to

cause an increase in ionisation energies.

IONISATION ENERGY

2) The distance of the electron from the nucleus

It has been found that, if F is the force of attraction between

two objects and d is the distance between them, then F is

proportional to1

𝑑2.

This distance effect means that all forces of attraction

decrease rapidly as the distance between the attracted bodies

increases. Thus the attraction between a nucleus and electrons

decrease as the quantum numbers of the shells increase.

The further the shell is from the nucleus, the lower are the

ionisation energies for electrons in that shell.

IONISATION ENERGY

3) The ‘shielding’ effect by electrons in filled inner shells

All electrons are negatively charged and repel each

other.

Electrons in the filled inner shells repel electrons in the

outer shell and reduce the effect of the positive nuclear

charge.

This is called the shielding effect. The greater the

shielding effect upon an electron, the lower is the energy

required to remove it and thus the lower the ionisation

energy.

IONISATION ENERGY

Hydrogen has an electronic structure of 1s1. It is a very small

atom, and the single electron is close to the nucleus and

therefore strongly attracted. There are no electrons screening it

from the nucleus and so the ionisation energy is high (1310 kJ

mol-1).

Helium has a structure 1s2. The electron is being removed from

the same orbital as in hydrogen's case. It is close to the nucleus

and unscreened. The value of the ionisation energy (2370 kJ

mol-1) is much higher than hydrogen, because the nucleus now

has 2 protons attracting the electrons instead of 1.

Lithium is 1s22s1. Its outer electron is in the second energy

level, much more distant from the nucleus. You might argue that

that would be offset by the additional proton in the nucleus, but

the electron doesn't feel the full pull of the nucleus - it is

screened by the 1s2 electrons. Lithium's first ionisation energy

drops to 519 kJ mol-1

PERIODIC PATTERN OF ∆Hi1

The first ionisation energies of the first 36 elements in the

Periodic Table are plotted against their atomic number and

presented in the following graph:


First Ionisation Energy Across a Period

From the graph, we see a general trend of increasing

ionisation energies across a Period. However, the trend is

uneven.

For example, at element 3 (lithium 1s22s1), 4 (beryllium

1s22s2) and 5 (boron 1s22s22p1). We might have predicted

boron would have the highest ionisation energy of the three;

in fact, it is beryllium.


Our modern theories for electronic structure show that the p

orbitals are higher energy level than the s orbital for a given

quantum number.

Hence we predict that an electron is more easily removed

from the p orbital than the s orbital. Thus, the 2p electron in

boron is easier to remove than one of the 2s electrons.

Therefore, though the nuclear charge in boron is larger than

in beryllium, boron has the lower first ionisation energy.


Let’s look at another example of elements in Period 2:

In the general trend across the Period, we might expect the

ionisation energy of oxygen to be higher than that of nitrogen.

In fact, the ionisation energy of nitrogen is the higher of the

two.

Nitrogen has three electrons in the p orbitals, each of them

unpaired; Oxygen has four electrons, with two of them paired.

The repulsion between the electrons in the pair increases the

energy and makes it easier to remove one of them and to ionise

an atom of oxygen, even though the nuclear charge is larger

than in an atom of nitrogen.


First Ionisation Energies in Groups

Elements are placed in Groups in the Periodic Table, as they

show many similar physical and chemical properties.

The first ionisation energies generally decrease down a

vertical Group, with increasing atomic number.

With increasing proton number, in any Group:

The positive nuclear charge increases;

The atomic radius increases so the distance of the outer

electrons from the nucleus also increases with each new

shell;

The shielding effect of the filled inner electron shells

increases as the number of inner shells grows.

The distance and shielding effects together reduce the effect

of the increasing nuclear charges from element to element

down any Group.



Ionisation Energy

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