Ionisation Energy Trends: GCSE Chemistry Explained

What exactly is first ionisation energy?

The first ionisation energy is the minimum energy needed to remove one electron from each atom in one mole of gaseous atoms, producing one mole of gaseous ions with a +1 charge. The process for element X is written as:

X(g) → X⁺(g) + e⁻ ΔH = first ionisation energy (kJ mol⁻¹)

Three conditions must be stated precisely:

• The atom must be in the gaseous state — removing an electron from a solid involves different energetic contributions.
• The ion formed carries a +1 charge (one electron removed).
• The process is endothermic — energy must be supplied, never released.

Ionisation energies are always positive values. Higher values mean the electron is held more tightly and harder to remove.

What three factors determine ionisation energy?

Three factors compete to determine how easily an electron is removed:

• Nuclear charge (proton number): A higher nuclear charge means the nucleus exerts a stronger attractive force on the outer electrons. More protons → higher ionisation energy, all else being equal.
• Atomic radius: The further an electron is from the nucleus, the weaker the electrostatic attraction (following Coulomb's law, force decreases with the square of distance). Larger atoms → lower ionisation energy.
• Electron shielding: Inner-shell electrons repel outer-shell electrons and partially cancel the pull of the nucleus. More inner shells → more shielding → lower effective nuclear charge → lower ionisation energy.

The term effective nuclear charge (Z_eff) captures the net pull experienced by an outer electron after shielding is accounted for. Z_eff = nuclear charge − shielding from inner electrons.

What is the trend in ionisation energy across Period 3?

Across Period 3 (sodium to argon), the first ionisation energy generally increases from left to right.

Reason for the general increase: Across a period, electrons are added to the same shell, so shielding from inner shells stays approximately constant. However, the nuclear charge increases by one proton with each element. The outer electrons experience a stronger effective nuclear charge and are held more tightly.

Why does Al have a lower first ionisation energy than Mg? Magnesium's outer electron is in a 3s orbital; aluminium's outer electron is in a 3p orbital. The 3p orbital is slightly higher in energy and is partially shielded by the 3s electrons, making it easier to remove despite aluminium having a greater nuclear charge.

Why does S have a lower first ionisation energy than P? Phosphorus has three electrons in separate 3p orbitals (Hund's rule — one per orbital), whereas sulfur has four 3p electrons, meaning one orbital holds a pair. Electron–electron repulsion within that paired orbital makes one electron easier to remove than the unpaired electrons in phosphorus.

What is the trend in ionisation energy down Group 1?

Down Group 1 (lithium to caesium), first ionisation energy decreases with each successive element.

Reason for the decrease: Each element down the group has one additional electron shell compared to the element above it. The outer electron is therefore further from the nucleus (larger atomic radius) and there are more inner shells providing shielding. The increased distance and increased shielding both reduce the effective nuclear charge experienced by the outer electron, making it progressively easier to remove. Although nuclear charge also increases down the group, the increase in shielding and atomic radius outweighs this effect.

How do ionisation energies link to the periodic table?

The successive ionisation energies of a single element reveal its electronic structure. When there is a large jump between consecutive ionisation energies, this marks the transition from removing outer-shell electrons to inner-shell electrons. For example, sodium shows a sharp jump after the first ionisation energy (removing the single outer electron) and another after the ninth (once all of the second shell is removed). These jumps directly confirm the shell structure of atoms.

The pattern of ionisation energies across periods also explains the reactivity trends you observe:

• Metals on the left of the periodic table have low ionisation energies → readily form positive ions → behave as reducing agents
• Non-metals on the right have high ionisation energies → do not easily lose electrons → tend to gain electrons instead

How is ionisation energy relevant to bonding?

Elements with low ionisation energies can easily donate electrons to form positive ions in ionic bonds (e.g., sodium, potassium). Elements with very high ionisation energies, such as chlorine or fluorine, find it energetically unfavourable to lose electrons and instead form covalent bonds by sharing electrons. Noble gases have the highest ionisation energies in each period — combined with their already-full outer shells, this explains their extreme chemical inertness.

Frequently asked questions

Q: What is the difference between first and second ionisation energy?

The first ionisation energy removes the first electron from a neutral gaseous atom (X → X⁺). The second ionisation energy removes a second electron from the +1 ion already formed (X⁺ → X²⁺). Second ionisation energies are always higher than first, because the same nuclear charge now acts on fewer electrons — the remaining electrons are held more tightly, and the removal of a second electron may also mean breaking into an inner shell.

Q: Why is ionisation energy measured in gaseous atoms?

If atoms were in a solid or liquid state, additional energy contributions from lattice energy or intermolecular forces would be included in any measurement, making comparisons between elements meaningless. Using gaseous atoms as the starting state isolates the pure nuclear–electron interaction and gives a consistent, comparable value.

Q: Does ionisation energy affect the colour of flame tests?

Flame tests are related to but distinct from ionisation energies. In a flame, electrons are promoted to higher energy levels (excited states) by thermal energy. When they return to the ground state, they emit light at wavelengths specific to the energy gap — producing the characteristic colours. While the energy levels involved in flame tests are related to the overall electronic structure that ionisation energies describe, flame test colours are determined by emission spectra rather than by the first ionisation energy directly.

Q: Why does ionisation energy increase so sharply at noble gases?

Noble gases (helium, neon, argon, etc.) have completely filled outer electron shells. Their electrons are held exceptionally tightly because the nuclear charge is high relative to the number of shielding inner electrons, and there is no partly filled subshell available. This maximally stable configuration means removing an electron requires far more energy than for the preceding halogen, producing a sharp peak at each noble gas in the Period 3 ionisation energy graph.

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