Ionisation Energy and Electronic Structure

Specification Reference Topic 1, points 11–15 (Edexcel A-Level Chemistry)

Quick Notes

First Ionisation Energy
- Energy needed to remove one mole of electrons from one mole of gaseous atoms
- X(g) → X⁺(g) + e⁻
Successive Ionisation Energies
- Energy needed to remove each further electron from the same atom
- Each one is higher due to increased attraction (fewer electrons, same nuclear charge)
Factors Affecting Ionisation Energy
- Nuclear charge: more protons = stronger attraction = higher IE
- Electron shielding: more inner shells = more shielding = lower IE
- Atomic radius: greater distance = weaker attraction = lower IE
Trends in Ionisation Energy
- Across a period: increases (more protons, similar shielding)
- Down a group: decreases (more shielding, larger radius)
Evidence for Electronic Structure
- Emission spectra: show discrete energy levels (shells)
- Successive IEs: large jumps indicate new electron shells
- First IE trends: exceptions to general increase across a period reveal sub-shell structure (e.g. Mg to Al)

Full Notes

First Ionisation Energy

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms, to form one mole of gaseous 1+ ions.

Equation:
X(g) → X⁺(g) + e⁻

Ionisation energy is measured in kJ mol⁻¹. It gives information about how strongly the outermost electrons are held by the atom.

Successive Ionisation Energies

After the first electron is removed, further ionisation energies can be measured:

Second ionisation energy:
X⁺(g) → X²⁺(g) + e⁻
Third ionisation energy:
X²⁺(g) → X³⁺(g) + e⁻

Each successive ionisation energy is larger than the one before because the electrons are being removed from a positively charged ion. The same nuclear charge is attracting fewer electrons and the remaining electrons are held more tightly.

Sharp increases between ionisation energy values indicate electrons are being removed from a new shell (see below).

Factors Affecting Ionisation Energy

Ionisation energy is influenced by three key factors:

Nuclear charge
More protons in nucleus = greater positive charge = stronger attraction to electrons = higher ionisation energy
Electron shielding
Inner electrons shield outer electrons from the nucleus = weaker attraction = lower ionisation energy
Atomic radius
Greater distance between outermost electron and the nucleus = weaker attraction = lower ionisation energy

These factors work together to determine how tightly electrons are held and how difficult they are to remove.

Trends in Ionisation Energy

Trends in ionisation energies help explain the reactivity of elements in different groups and periods in the periodic table.

Across a Period:

Ionisation energy increases as a trend

Reason: More protons in nucleus = greater nuclear charge. Electrons added to the same shell, so shielding (repulsion from inner electrons) remains similar, but attraction to the nucleus increases.

For example across period 3

Edexcel A-Level Chemistry line graph showing first ionisation energy increasing across a period with noted small dips indicating sub-shell effects.

Down a Group:

Ionisation energy decreases

Reason: Atomic radius increases meaning more shielding from inner electrons. This causes weaker attraction between outer electrons and nucleus, meaning outer electron is easier to remove.

For example down group 2

Edexcel A-Level Chemistry plot showing first ionisation energy decreasing down a group due to increased shielding and larger atomic radius.

Evidence for Electronic Structure

Several types of data support the current model of electronic structure and configurations:

Atomic emission spectra

Shows that electrons exist in quantised energy levels (shells)
When electrons drop from higher to lower energy levels, they emit energy at specific wavelength, showing that they can only have specific amounts of energy (only exist in specific energy levels).

For example The hydrogen emission spectra shows the specific energies emitted by electrons as they fall to the n = 2 energy level from higher energy levels in a hydrogen atom.

Edexcel A-Level Chemistry diagram of hydrogen emission spectrum illustrating discrete lines corresponding to electron transitions to n = 2.

Successive ionisation energies

Sharp increases show when an inner shell is reached
The number of electrons removed before the large jump indicates the number of electrons in the outer shell

Exceptions to trends in first ionisation energies

Exceptions to the general increase in first ionisation energy across periods 2 and 3 support the existence of sub-shells (e.g. s and p)

Edexcel A-Level Chemistry plot showing first ionisation energy across periods with dips at Al and S evidencing sub-shell structure and electron pairing.

Aluminium (Al): Lower than magnesium (Mg), meaning electron removed must have less attraction to the nucleus (be further away and higher energy). This is now explained as the outer electron being in a 3p orbital (higher in energy than the 3s orbital the outer electron in magnesium is in).

Sulfur (S): Lower than phosphorus (P), meaning outermost electron removed must be higher in energy and easier to remove than for sulfur. This is now explained as being due to electron pairing in a 3p orbital, causing repulsion and giving evidence that the 3p sub shell must contain 3 orbitals (as the 4th electron in the sub shell has to pair up with another electron in an orbital).

Together, these observations confirm the shell and sub-shell model of the atom.

Summary

First ionisation energy removes one mole of electrons from one mole of gaseous atoms.
Successive ionisation energies increase because the ion becomes more positive.
Nuclear charge increases IE, shielding and distance decrease IE.
Across a period IE increases and down a group IE decreases.
Emission spectra, successive IE jumps and first IE exceptions support the idea of shells and sub shells.